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1-1-11
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
1-1-22
11
Covalent Covalent Bonding & Bonding & Shapes of Shapes of MoleculesMolecules
Chapter 1Chapter 1
1-1-33
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
1-1-44
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
1-1-55
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
1-1-66
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
1-1-77
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
1-1-88
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
1-1-99
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 :
:::
::
.
.
.
.
.
.
.
..
..
. .
.
.
.
:
:
:
::::::::
::::::.
:::
:
1-1-1010
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
1-1-1111
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
1-1-1212
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-
1-1-1313
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--
1-1-1414
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
1-1-1515
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
::
:
1-1-1616
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
1-1-1717
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
1-1-1818
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
+
1-1-1919
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-
1-1-2020
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
: :
: :
: :
::
::
:
:
:
::
1-1-2121
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
:
:
:
::
::
::
::::
::
:
:
::
:
:
1-1-2222
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
1-1-2323
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
1-1-2424
11 AlcoholAlcohol contains an -OH (hydroxyl) group
H-C-C-O-H
H
H
H
H
Ethanol
(an alcohol)
::
1-1-2525
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)
: : :
1-1-2626
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
: :: :
1-1-2727
11 Carboxylic AcidCarboxylic Acid contains a carboxyl (-COOH) group
or orCH3-C-O-H CH3COOH CH3CO2H
Acetic acid(a carboxylic acid)
O: :::
1-1-2828
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
1-1-2929
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
:
1-1-3030
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-
1-1-3131
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
1-1-3232
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
1-1-3333
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
1-1-3434
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
1-1-3535
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
1-1-3636
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: :
::
1-1-3737
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!
1-1-3838
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
1-1-3939
11 ResonanceResonance Examples of ions and a molecule best
represented as resonance hybrids
carbonate ionCO32-
acetate ion
CH3COCH3acetone
nitrate ion
CH3COO-
NO3-
1-1-4040
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
1-1-4141
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
1-1-4242
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
1-1-4545
11 Shapes of 1s and 2s OrbitalsShapes of 1s and 2s Orbitals
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1-1-4646
11 Shapes of a Set of 2p Atomic OrbitalsShapes of a Set of 2p Atomic Orbitals
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1-1-4747
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
1-1-4848
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
1-1-4949
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
1-1-5252
11 Molecular OrbitalsMolecular Orbitals pi bonding and antibonding MOs
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1-1-5353
11 Molecular OrbitalsMolecular Orbitals computed pi bonding and antibonding MOs for
ethylene
1-1-5454
11 Molecular OrbitalsMolecular Orbitals computed pi bonding and antibonding orbitals
for formaldehyde
1-1-5555
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
1-1-5656
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
1-1-5757
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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1-1-5858
11 Bonding in CHBonding in CH44, NH, NH33, and H, and H22OO
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1-1-5959
11 spsp22 Hybrid Orbitals Hybrid Orbitals
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1-1-6060
11 Bonding in CHBonding in CH22=CH=CH22
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1-1-6161
11 Bonding in CHBonding in CH22=O=O
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1-1-6262
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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1-1-6363
11Bonding in CBonding in C22HH22
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1-1-6464
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