1-1 1 Organic Chemistry William H. Brown & Christopher S. Foote.

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11

Organic Organic ChemistryChemistry

William H. Brown & William H. Brown & Christopher S. FooteChristopher S. FooteWilliam H. Brown & William H. Brown & Christopher S. FooteChristopher S. Foote

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11

Covalent Covalent Bonding & Bonding & Shapes of Shapes of MoleculesMolecules

Chapter 1Chapter 1

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11 Organic ChemistryOrganic Chemistry The study of the compounds of carbon Over 10 million structures have been identified• about 1000 new ones are identified each day!

C is a small atom • it forms single, double, and triple bonds• it is intermediate in electronegativity (2.5)• it forms strong bonds with C, H, O, N, and some metals

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11 Structure of Atoms ElectronicStructure of Atoms Electronic Structure of atoms• small dense nucleus,

diameter 10-14 - 10-15 m, which contains most of the mass of the atom

• extranuclear space, diameter 10-10 m, which contains positively-charged electrons

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11 Electrons are confined to regions of space

called principle energy levels (shells)• each shell can hold 2n2 electrons (n = 1,2,3,4......)

Structure of Atoms ElectronicStructure of Atoms Electronic

Shell

Number of Electrons ShellCan Hold

Relative Energiesof Electrons in These Shells

3218 8 2

4321

higher

lower

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11 Electronic Structure of AtomsElectronic Structure of Atoms Shells are divided into subshells called orbitals, which are

designated by the letters s, p, d, f,........• s (one per shell)

• p (set of three per shell 2 and higher)

• d (set of five per shell 3 and higher) .....

Shell Orbitals Contained in That Shell

3

2

1 1s

2s, 2px, 2py, 2pz

3s, 3px, 3py, 3pz, plus five 3d orbitals

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11 Electron Structure of AtomsElectron Structure of Atoms Aufbau Principle:Aufbau Principle: orbitals fill from lowest to

highest energy Pauli Exclusion Principle:Pauli Exclusion Principle: only two electrons per

orbital, spins must be paired Hund’s Rule:Hund’s Rule: for a set of degenerate orbitals, add

one electron in each before a second is added in any one

Example:Example: Write the ground-state electron configuration for each element

(a) Li (b) O (c) Cl

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11 Lewis StructuresLewis Structures Gilbert N. Lewis Valence shell:Valence shell: the outermost electron shell of an

atom Valence electrons:Valence electrons: electrons in the valence shell

of an atom; these electrons are used to form chemical bonds

Lewis structure:Lewis structure: • the symbol of the atom represents the nucleus and all

inner shell electrons• dots represent valence electrons

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11 Lewis StructuresLewis Structures Table 1.4 Lewis Structures for Elements 1-18

N OB

H

Li Be

Na

He

Cl

F

S

Ne

Ar

C

SiAl P

1A 2A 3A 4A 5A 6A 7A 8A

Mg :

:::

::

.

.

.

.

.

.

.

..

..

. .

.

.

.

:

:

:

::::::::

::::::.

:::

:

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11 Formation of Chemical BondsFormation of Chemical Bonds Atoms bond together so that each atom acquires

an electron configuration the same as the noble gas nearest it in atomic number• an atom that gains electrons becomes an anion• an atom that loses electrons becomes a cation

Two extremes of bonding• ionic bond:ionic bond: a chemical bond resulting from the

electrostatic attraction of an anion and a cation• covalent bond:covalent bond: a chemical bond formed between two

atoms by sharing one or more pairs of electrons

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11 ElectronegativityElectronegativity Electronegativity:Electronegativity: a measure of the force of an atom’s attraction

for the electrons it shares with another atom in a chemical bond Pauling scale

• increases left to right in a row

• increases bottom to top in a column

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11 ElectronegativityElectronegativityTable 1.6 Classification of Chemical Bonds

Difference in ElectronegativityBetween Bonded AtomsType of Bond

Less than 0.5

0.5 to 1.9

Greater than 1.9

Nonpolar covalentPolar covalent

Ionic

H Cl2.1 3.0

δ+ δ-

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11 Bond DipolesBond Dipoles Table 1.7 Average Bond Dipoles of Selected Covalent

Bonds

C-FC-ClC-BrC-I

H-OH-NH-C C-O

C-NH-S

C=O

Bond

1.41.51.41.2

Bond

Bond Dipole (D)

1.51.30.3 0.7

0.20.7

2.3

3.5

Bond Dipole (D)

Bond Dipole (D) Bond

C=N--

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11 Lewis StructuresLewis Structures To write a Lewis structure• determine the number of valence electrons• determine the arrangement of atoms• connect the atoms by single bonds• arrange the remaining electrons so that each atom has

a complete valence shell• show bonding electrons as a single bond (a single

line); show nonbonding electrons as a pair of dots• in a single bond atoms share one pair of electrons, in a

double bond they share two pairs of electrons, and in a triple bond they share three pairs of electrons

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11 Lewis StructuresLewis Structures

:::H O

H

H

H NH C

H

H

H Cl

H

H

H2O (8)

NH3 (8)CH4 (8)

HCl (8)

Hydrogen chloride

Methane Ammonia

Water

::

:

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11 Lewis StructuresLewis Structures In neutral molecules• hydrogen has one bond• carbon has 4 bonds and no unshared electrons• nitrogen has 3 bonds and 1 unshared pair of electrons• oxygen has 2 bonds and 2 unshared pairs of electrons• halogens have 1 bond and 3 unshared pairs of

electrons

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11 Formal ChargeFormal Charge Formal charge:Formal charge: the charge on an atom in a

molecule or polyatomic ion To derive formal charge

1. write a correct Lewis structure for the molecule or ion

2. assign each atom all of its unshared (nonbonding) electrons and one-half its shared (bonding) electrons

3. compare this number with the number of valence electrons in the neutral, unbonded atom

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11 Formal ChargeFormal Charge

• if the number assigned to the bonded atom is less than that assigned to the unbonded atom, the atom has a positive formal charge• if the number is greater, the atom has a negative formal charge

number of valence electrons in the neutral, unbonded atom

all unsharedelectrons

one half of all shared electrons

+

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11 Formal ChargeFormal Charge

Example:Example: Draw Lewis structures and show all formal charges for these ions

(a) (b) (c)NH2- HCO3

- CO32-

(d) (e) (f)NO3- HCOO- CH3COO-

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11 Exceptions to the Octet RuleExceptions to the Octet Rule Molecules containing atoms of Group 3A

elements, particularly boron and aluminum

::

:

F B

F

F

Cl Al

Cl

Cl

6 electrons in the valence shells of boron

and aluminum

Boron trifluoride Aluminum chloride

: :

: :

: :

::

::

:

:

:

::

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11 Exceptions to the Octet RuleExceptions to the Octet Rule Atoms of third-period elements have 3d orbitals

and may expand their valence shells to contain more than 8 electrons• phosphorus may have up to 10

:

:

Phosphorus pentachloride

Phosphoricacid

P

ClCl Cl

Cl ClCH3-P-CH3

CH3

Trimethyl-phosphine

H-O-P-O-H

O

O-H

:

:

:

::

::

::

::::

::

:

:

::

:

:

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11• and S, which may have up to 12 electrons in its

valence shell

:

::H-S-H CH3-S-CH3 H-O-S-O-H

O

OO

Sulfuricacid

Hydrogensulfide

Dimethylsulfoxide

::

::

: : : :

: :

Exceptions to the Octet RuleExceptions to the Octet Rule

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11 Functional GroupsFunctional Groups Functional group:Functional group: an atom or group of atoms

within a molecule that shows a characteristic set of physical and chemical properties

Functional groups are important for three reason; they are1. the units by which we divide organic compounds into

classes

2. the sites of characteristic chemical reactions

3. the basis for naming organic compounds

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11 AlcoholAlcohol contains an -OH (hydroxyl) group

H-C-C-O-H

H

H

H

H

Ethanol

(an alcohol)

::

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11 AmineAmine contains an amino group; a nitrogen bonded to

one, two, or three carbon atoms• may by 1°, 2°, or 3°

CH3 N H

H

CH3 N H

CH3

CH3 N CH3

CH3

Methylamine(a 1° amine)

Dimethylamine(a 2° amine)

Trimethylamine(a 3° amine)

: : :

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11 Aldehyde and KetoneAldehyde and Ketone contains a carbonyl (C=O) group

Acetaldehyde(an aldehyde)

Acetone(a ketone)

O O

CH3-C-H CH3-C-CH3

: :: :

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11 Carboxylic AcidCarboxylic Acid contains a carboxyl (-COOH) group

or orCH3-C-O-H CH3COOH CH3CO2H

Acetic acid(a carboxylic acid)

O: :::

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11 Carboxylic EsterCarboxylic Ester a derivative of a carboxylic acid in which the

carboxyl hydrogen is replaced by a carbon group

CH3-C-O-CH2-CH3

Ethyl acetate(An ester)

::

: :O

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11 VSEPR ModelVSEPR Model

HC C

H

O C

C

H

NH

H

C

HH

OH

CH C H

H

O

4 regions of e- density(tetrahedral, 109.5°)

3 regions of e- density(trigonal planar, 120°)

2 regions of e- density(linear, 180°)

H

CH H

HN

H HH

::

::

::

::

:

H OH

:

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11 VSEPR ModelVSEPR Model Example:Example: predict all bond angles for these molecules and

ions

(a) NH4+ (b) CH3NH2

(f) H2CO3 (g) HCO3-(e) CH3CH=CH2

(i) CH3COOH(h) CH3CHO

(d) CH3OH

(j) BF4-

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11 To determine if a molecule is polar, we need to

determine • if the molecule has polar bonds• the arrangement of these bonds in space

Dipole moment (Dipole moment ():): the vector sum of its individual bond dipole moments in a molecule• reported in debyes (D)

Polar and Nonpolar MoleculesPolar and Nonpolar Molecules

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11 these molecules have polar bonds, but each has

a zero dipole moment

O C O

Carbon dioxide = 0 D

B

F

F

F

Boron trifluoride = 0 D

C

Cl

ClClCl

Carbon tetrachloride = 0 D

Polar and Nonpolar MoleculesPolar and Nonpolar Molecules

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11 These molecules have polar bonds and a dipole

moment greater than zero

Polar and Nonpolar MoleculesPolar and Nonpolar Molecules

N

HH

H

OH H

Water = 1.85D

Ammonia = 1.47D

direction of dipole

moment in ammonia

direction of dipole

moment in water H

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11 ResonanceResonance For many molecules and ions, no single Lewis

structure provides a truly accurate representation

Ethanoate ion(Acetate ion)

C

O

O

CH3C

O

O

CH3

-

-

and

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11 ResonanceResonance Linus Pauling - 1930s• many molecules and ions are best described by

writing two or more Lewis structures• individual Lewis structures are called contributing

structures• connect individual contributing structures by double-

headed (resonance) arrows• the molecule or ion is a hybrid of the various

contributing structures

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11 ResonanceResonance Examples:Examples: equivalent contributing structuresequivalent contributing structures

Acetate ionNitrite ion:

:

:

:: : :

N

O -

O

: N

O

O -

:

: : :

:

:

: :

: :C

O -

OH3C C

O

O -

H3C: :

::

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11 ResonanceResonance Curved arrow:Curved arrow: a symbol used to show the

redistribution of valence electrons In using curved arrows, there are only two

allowed types of electron redistribution:• from a bond to an adjacent atom• from an atom to an adjacent bond

Electron pushing is a survival skill in organic chemistry. Learn it well!

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11 All contributing structures must1. have the same number of valence electrons

2. obey the rules of covalent bonding• no more than 2 electrons in the valence shell of H • no more than 8 electrons in the valence shell of a 2nd

period element• a 3rd period element may have up to 12 electrons in

its valence shell

3. differ only in distribution of valence electrons

4. have the same number of paired and unpaired electrons

ResonanceResonance

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11 ResonanceResonance Examples of ions and a molecule best

represented as resonance hybrids

carbonate ionCO32-

acetate ion

CH3COCH3acetone

nitrate ion

CH3COO-

NO3-

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11 ResonanceResonance Preference 1:Preference 1: filled valence shells• structures in which all atoms have filled valence shells

contribute more than those with unfilled valence shells

••

••••

Greater contribution; both carbon and oxygen have complete valence shells

Lesser contribution;carbon has only 6 electrons in its valence shell

+ +CCH3 OCH3 O

H

HC

H

H

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11 ResonanceResonance Preference 2:Preference 2: maximum number of covalent

bonds• structures with a greater number of covalent bonds

contribute more than those with fewer covalent bonds

••

••••

Greater contribution(8 covalent bonds)

Lesser contribution(7 covalent bonds)

+ +CCH3 OCH3 O

H

HC

H

H

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11 ResonanceResonance Preference 3:Preference 3: least separation of unlike charge• structures with separation of unlike charges

contribute less than those with no charge separation

CH3-C-CH3 CH3-C-CH3

Greater contribution(no separation of unlike charges)

Lesser contribution(separation of unlike

charges)

O -O::

:::

1-1-4343

11 ResonanceResonance Preference 4:Preference 4: negative charge on the more

electronegative atom • structures that carry a negative charge on the more

electronegative atom contribute more than those with the negative charge on the less electronegative atom

CH3CH3H3CCH3H3CC+

O -

C

O

H3CC-

O +

Greater contribution

Can be ignored Lessercontribution

: : :

:

:: : :

1-1-4444

11 Quantum or Wave MechanicsQuantum or Wave Mechanics Albert Einstein: E=h (energy is quantized)• light has particle properties

Louis deBroglie: wave/particle duality Erwin Schrödinger: wave equation• 2 is the probability of finding an electron in a given

region of space • the standard representation of an atomic orbital is a

boundary surface representing 95% probability of finding an electron in that region of space

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11 Shapes of 1s and 2s OrbitalsShapes of 1s and 2s Orbitals

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11 Shapes of a Set of 2p Atomic OrbitalsShapes of a Set of 2p Atomic Orbitals

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11 Molecular Orbital TheoryMolecular Orbital Theory• electrons in atoms exist in atomic orbitals• electrons in molecules exist in molecular orbitals

(MOs)• using the Schrödinger equation, we can calculate the

shapes and energies of MOs

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11 Molecular Orbital TheoryMolecular Orbital Theory Rules:• combination of n atomic orbitals gives n MO• MOs are arranged in order of increasing energy• MOs filling is governed by the same rules as for

atomic orbitals:• Aufbau principle: fill beginning with LUMO• Pauli exclusion principle: no more than 2e- in a MO• Hund’s rule: filling of degenerate orbitals

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11 Molecular Orbital TheoryMolecular Orbital Theory Terminology• ground state = lowest energy• excited state = NOT lowest energy• = sigma bonding MO• * = sigma antibonding MO• = pi bonding MO• * = pi antibonding MO

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11 Molecular Orbital TheoryMolecular Orbital Theory sigma 1s bonding and antibonding MOs

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11 Molecular OrbitalsMolecular Orbitals computed sigma bonding and antibonding MOs

for H2

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11 Molecular OrbitalsMolecular Orbitals pi bonding and antibonding MOs

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11 Molecular OrbitalsMolecular Orbitals computed pi bonding and antibonding MOs for

ethylene

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11 Molecular OrbitalsMolecular Orbitals computed pi bonding and antibonding orbitals

for formaldehyde

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11 Hybrid OrbitalsHybrid Orbitals The Problem:• bonding by 2s and 2p atomic orbitals would give bond

angles of approximately 90°• instead we observe bond angles of approximately

109.5°, 120°, and 180°

A Solution• hybridization of atomic orbitals• 2nd row elements use sp3, sp2, and sp hybrid orbitals

for bonding

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11 Hybrid OrbitalsHybrid Orbitals Hybridization of orbitals (L. Pauling)• the combination of two or more atomic orbitals forms a

new set of atomic orbitals, called hybrid orbitals

We deal with three types of hybrid orbitalsspsp33 (one s orbital + three p orbitals)

spsp22 (one s orbital + two p orbitals)

spsp (one s orbital + one p orbital)

Overlap of hybrid orbitals can form two types of bonds depending on the geometry of overlap bondsbonds are formed by “direct” overlap

bondsbonds are formed by “parallel” overlap

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11 spsp33 Hybrid Orbitals Hybrid Orbitals• each sp3 hybrid orbital has

two lobes of unequal size• the sign of the wave

function is positive in one lobe, negative in the other, and zero at the nucleus

• the four sp3 hybrid orbitals are directed toward the corners of a regular tetrahedron at angles of 109.5°

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11 Bonding in CHBonding in CH44, NH, NH33, and H, and H22OO

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11 spsp22 Hybrid Orbitals Hybrid Orbitals

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11 Bonding in CHBonding in CH22=CH=CH22

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11 Bonding in CHBonding in CH22=O=O

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11 sp Hybrid Orbitalssp Hybrid Orbitals• two lobes of unequal size at an angle of 180°• the two unhybridized 2p orbitals are perpendicular to

each other and to the line through the two sp hybrid orbitals

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11Bonding in CBonding in C22HH22

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11 Hybrid OrbitalsHybrid Orbitals

H-C C-H

C C

H-C-C-H

OrbitalHybrid-ization

Types of Bonds to Carbon Example

sp3 four sigma bonds

sp2 three sigma bondsand one pi bond

sp two sigma bondsand two pi bonds

Ethane

Ethylene

Acetylene

Name

PredictedBondAngles

109.5°

120°

180°

H

H

H

H

H

HH

H

1-1-6565

11 Hybrid OrbitalsHybrid Orbitals

H-C-C-H

H

H

H

H

HC C

H

H H

H-C C-H

C-C

C-C

C-C

C-H

C-H

C-H

FormulaBondOrbitalOverlap

Bond Length(pm)

Bond Strength[kJ (kcal)/mol]

sp3-sp3

sp2-1s

sp-sp, two 2p-2p

sp-1s

sp3-1s

sp2-sp2, 2p-2p

153.2

111.4

133.9

110.0

121.2

109.0

368 (88)

410 (98)

611 (146)

435 (104)

837 (200)

523 (125)

Ethane

Ethylene

Acetylene

Name

1-1-6666

11 Prob 1.26Prob 1.26Write Lewis structures for these molecules

(a) (b) (c)H2O2 N2H4 CH3OH

(d) (e) (f)CH3SH CH3NH2 CH2Cl2

(g) (h) (i)CH3OCH3 N2H4 CH3OH

(j) (k) (l)CH3COOH CH3COCH3 HCN

(m) (n) (o)HNO3 HNO2 HCOOH

1-1-6767

11 Prob 1.27Prob 1.27Write Lewis structures for these ions

(a) (b) (c)NH2- HCO3

- CO32-

(d) (e) (f)NO3- HCOO- CH3COO-

1-1-6868

11 Prob 1.28Prob 1.28 Complete the structural formula and write the molecular

formula of each compound.

(a)C-C=C-C-C

C

(b) (c)

(d) (e) (f)

(g) (h) (i)

C-C-C-H

C-C-C-C-C

C-C-C-C-NH2

C-C-C-C-OH C=C-C-OH

C-C-C-OH

C-C-C-C-OH C-C-C-CO O

C

O

OH OOH

NH2

O

C

C

1-1-6969

11 Prob 1.29Prob 1.29 Which structural formulas are incorrect, and why?

H-C-C=O-HH

H H

H(a) (b) (c)H-C=C-H H-N-C-C-O-H

H

Cl

H H

H

H H

H

H-C C-C-HH

H

H(d) (e) (f)H-O-C-C-C-O-H H-C-C-C-H

O

H

H

H

H

H

H

H

H

H

O

(g) (h)H-C-C=C=C-C-HH

H

H H H

HC CH C-H

H

H

1-1-7070

11 Prob 1.30Prob 1.30 Add valence electrons to complete the outer shell of each

atom and assign any formal charges

(a) C OCHH

H

H

H

(b) (c) C CNHCH

HCH

HH

H

H

O

O

H

H

1-1-7171

11 Prob 1.31Prob 1.31Assign all formal charges

H-C-C-C-H

H

H

H

O

:

::

(a) (b) (c)

(d) (e) (f)H- C-C= C-H H- C-C- C- H H- C-O- H

H- C-C- H

H

O

H

H

H

O

H

H

H H

H

H

H

H

::

::

:

:

H- N- C= C-H

H H

: O:::

1-1-7272

11 Prob 1.37Prob 1.37 Use the VSEPR model to predict all bond angles

:(a) (b) (c)

(d) (e) (f)

H-C-C-O-H

H

H H

H

H-C=C-Cl H-C-C

H-C-N-HH-O-N=OH-C-O-H

H HC-H

H

H

O

H

H

::

::

::::

::

:

:H

: :

1-1-7373

11 Prob 1.38Prob 1.38 Use the VSEPR model to predict all bond angles

(a) (b) (c)

(d) (e) (f)

CH3-CH=CH2 CH3-N-CH3 CH3-CH2-C-OH

CH3-CH=N-OHCH2=C=OCH2=C=CH2

CH3 O

1-1-7474

11 Prob 1.39Prob 1.39 Use the VSEPR model to predict the geometry of each ion.

(a) (b) (c)

(d) (e) (f)

NH2- NO2

- NO2+

CH3-CH3COO-NO3

-

(g) AlCl4-

1-1-7575

11 Prob 1.47Prob 1.47Identify the functional groups in each compound.

CH3-CH-C-OHOH O

HO-CH2-CH2-OH

CH3-CH-C-OHO

NH2

HO-CH2-CH-C-HOOH

H2NCH2CH2CH2CH2CH2CH2NH2CH3-C-CH2-C-OHO O

(a) (b)

(c) (d)

(e) (f)

1-1-7676

11 Prob 1.48Prob 1.48 Which compounds have dipole moments greater than

zero? In what direction does it point?

(a) (c)(b)CH3F CH2Cl2 CH2ClBr

(d) (e) (f)CFCl3 CCl4 CH2=CCl2

(g) (h) (i)CH2=CHCl HC C C CH CH3C N

(j) (k)BrCH=CHBr (two answers)(CH3)2C=O

1-1-7777

11 Prob 1.55Prob 1.55 State the orbital hybridization of each highlighted atom.

(a) (b) (c) CCC CH

HH

HHH

(d) (e) (f)OCH

H

HC

H

OO H O HCH H

(g) (h) O

H

H

C

H

H N OH N H (i)CH2=C=CH2

: :

:

:

::

: : : :: :

(a) (b) (c) CCC CH

HHH

(d) (e) (f)OC C

H

OO H O HCH H

(h) O

H

C H N OH N H (i)CH2=C=CH2

:

:

:

::

: : : :: :

:

HC

H

H

HCH H

HCH

CH

:

1-1-7878

11 Prob 1.56Prob 1.56 Describe each highlighted bond in terms of the overlap of

atomic orbitals.

:

(b)(a) C HCHC C (c)H

H H

H

(f)(e)(d) OCHCH

H O HH

H

(i)(h)(g) OC O

O

H

H

HC

H H

H

H N OH CH N H

CH2=C=CH2

:

:

: ::

: :

:

:

: ::

:

O

C-O-HH

::

::

1-1-7979

11 Prob 1.57Prob 1.57 Predict all bond angles, the hybridization of each carbon,

and the shape of a benzene molecule.

CC

CC

C

C

H

H

H

H

H

H

1-1-8080

11 Prob 1.58Prob 1.58 Complete the Lewis structure and describe each

highlighted bond in terms of the overlap of atomic orbitals.

NC

S CH

CCH2-S-CH2-CH2-CC=N

H2N

H2N NH2

N-S-NH2

O

O

1-1-8181

11 Prob 1.61Prob 1.61 Draw a Lewis structure for each compound. Show

covalent bonds by dashes, and ionic bonds by the charge on each ion.

(a) (b) (c)

(d) (e)

CH3ONa NH4Cl NaHCO3

NaBH4 LiAlH4

1-1-8282

11 Prob 1.70Prob 1.70 Identify the atoms in the organic starting material that

change hybridization upon reaction, and what the change is.

(a)

(b)

HC C

H

H H+ Cl2

C C HH + Cl2H

C CCl

Cl H

H-C-C-HCl

H Cl

H

(c) + H2OC C HH H-C-C-HH

H O

1-1-8383

11 Prob 1.70Prob 1.70 Identify the atoms in the organic starting material that

change hybridization upon reaction, and what the change is.

(d) + H2H-C-H H-C-O-H

H

H

(e) + H2OH-C-C-C-H

H

H

H H

H

H-C-C-C-H

H

H

H H

HOH

+ H+

(f)H

C CH

H H+H-C-C-O-C-C-H

H

H

H

H

H

HOH-O-C-C-H

H

H

O

O

1-1-8484

11

Covalent Covalent Bonds & Bonds & Shapes of Shapes of MoleculesMolecules

End Chapter 1End Chapter 1