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Ch.2 Polar Bonds and Their Consequences
polar covalent bonds: electron distribution is unsymmetrical
Ionic Character
X Y X Y X Yδ+ δ- + -
symmetricalcovalent bond
polarcovalent bond ionic bond
2.1 Polar Covalent Bonds and Electronegativity
Ch.2 Polar Bonds and Their Consequences
- EN difference: 0.3-2.0 (polar bond), >2.0 (ionic)
Elecrtonegativity (EN): the intrinsic ability of an atom to attract electrons in a covalent bond
H C N O F
PSi S Cl
Br
I
4.02.52.1 3.53.0
3.01.8 2.52.1
2.8
2.5
Li
Na
K
1.0
0.9
0.8
Be
Mg1.6
1.2Ca1.0
Cs0.7
Ch.2 Polar Bonds and Their Consequences
δ-Cl
C
HH H
δ+ δ-
MgBr
C
HH H
δ+direction
of polarity
organometalliccompounds
Ch.2 Polar Bonds and Their Consequences
inductive effects: shifting of electrons in σ-bond in response of the electronegativity of nearby atoms (atom's ability to polarize a bond); polarize bonds, chemical reactivity
Cl
C
HH H
electronegative atoms inductively
withdraw electrons
metals inductivelydonates electrons
MgBr
C
HH H
Ch.2 Polar Bonds and Their Consequences
dipole moment (μ): net molecular polarity (vector sum); resulted from individual bond polarities and lone pair contributions; polar substances dissolve in polar solvents, nonpolar substances soluble in nonpolar solvents
µ = Q x r Q: charge at either end of dipoler: distance between the charges
1 debyes (D): 3.336 x 10-30 C . m
For one positive charge and one negative charge seperated by 100 pm
-B
A+100 pm
µ = Q x r = (1.60 x 10-19 C) x (100 x 10-12 m)(1D / 3.336 x 10-30 C.m)= 4.80 D
2.2 Polar Covalent Bonds and Dipole Moment
Ch.2 Polar Bonds and Their Consequences
δ-Cl
C
HH H
δ+
µ = 1.87 D
C-Cl: 178 pm
µ = 4. 80 D x 1.78 = 8.54 D
For fully seperated chloromethane (if the C-Cl bond were ionic, C+ Cl-)
but the actual dipole moment is 1.87 D
x 100= 22%
therefore, the chlorine atom in chloromethane has an excess of about 0.2 electron,and the carbon has a deficiency of about 0.2 electron
1.87
8.54
δ-Cl
C
HH H
δ+
0.2 electron excess
0.2 electron deficient
Ch.2 Polar Bonds and Their Consequences
nonbonding electron pair: stick out from positively charged nuclei
H
OH
netnetN
HH H
H
OH
N
HH H
Ch.2 Polar Bonds and Their Consequences
NaCl 9.0
Dipole moment(D)
H3C NO
O3.46
CH3ClH2OCH3OH
H2C=N=N
1.871.851.701.50
NH3 9.0
CH4
CCl4CH3CH3
000
0BF3
0
Dipole moment(D)
Table 2.1 dipole moments of some compounds
Ch.2 Polar Bonds and Their Consequences
• related to bond polarity, dipole moment; common for atoms that have an apparently abnormal number of bonds
C NO
O
H
H
H
formal positive chargeformal negative charge
Nitromethane
2.3 Formal Charges
Ch.2 Polar Bonds and Their Consequences
C C
H
HHH
4 valence electronsfor isolated carbon atom
N N HHH
5 valence electronsfor isolated nitrogen atom
electron counts- covalent bond: each atom owns one electron; single bond (1 e-); double bond (2 e-); triple bond (3 e-)
- non-bonding electrons: owned by the atom
8
2= 4 valence electrons
6
2+ 2 = 5 valence electrons
Ch.2 Polar Bonds and Their Consequences
Formal Charge = # of valence electronsin free atom
# of valence electronsin bound atom
_
# of valenceelectrons
half ofbondingelectrons
# ofnonbonding
electrons
_ _=
Ch.2 Polar Bonds and Their Consequences
C NO
O
H
H
H
Nitromethane
nitrogen valence electronsnitrogen bonding electronsnitrogen nonbonding electrons
= 5= 8= 0
oxygen valence electronsoxygen bonding electronsoxygen nonbonding electrons
= 6= 2= 6
For neutral molecules: sum of the formal charges equal to zero
Formal Charge = 5 - 8/2 - 0 = + 1
Formal Charge = 6 - 2/2 - 6 = - 1
Ch.2 Polar Bonds and Their Consequences
C C C
+ 1 0 - 1
Table 2.2 A summary of formal charges on atoms
• dipolar molecule: neutral overall but have +/- charges on individual atoms; dipolar character of molecules → chemical reactivity
N N N
+ 1 0 - 1
O O O
+ 1 0 - 1
H3CS
CH3
O
Practice formal charge
Dimethyl sulfoxide
H3CS
CH3
O
H3CS
CH3
Oor
H2C N N
Diazomethane
H2C N N
H3C C N
Acetonitrile
H3C C N
H3C N C
Methyl isocyanide
H3C N C
Ch.2 Polar Bonds and Their Consequences
Ch.2 Polar Bonds and Their Consequences
• For some molecules simple Lewis structure cannot describe the actual structure correctly
C NO
O
H
H
H
Nitromethane
C NO
O
H
H
H
-
2.4 Resonance
For example, nitromethane has two equivalent N-O bonds, but one Lewis structure can't represent it
Both oxygen-nitrogen bonds are 122 pm in length, midway between the lengthOf a typical N-O single bond (130 pm) and a typical N=O double bond (116 pm)
Ch.2 Polar Bonds and Their Consequences
resonance form: individual Lewis structures that represent one molecule
Ch.2 Polar Bonds and Their Consequences
1. Individual resonance forms are imaginary, not real; The real structure is composite (or resonance hybrid) of different forms.
C NO
O
H
H
H
Nitromethane
C NO
O
H
H
H
-
2.5 Rules for Resonance Forms
2. Resonance forms differ only in the placement of their π or nonbonding electrons; Neither the position nor the hybridization of any atom changes from one resonance form to another.
A curved arrow shows movement of two electrons.
Ch.2 Polar Bonds and Their Consequences
CC
CCC
CH
H
HH
H
H
CC
CCC
CH
H
HH
H
H
Benzene (two resonance forms)
Ch.2 Polar Bonds and Their Consequences
H3C CH3
O strongbase
H3C C
O
H
H H3C C
O
H
H
3. Different resonance forms of a substance don't have to be equivalent; The real structure is composite (or resonance hybrid) of different forms.
For non-equivalent resonance forms, the actual structure of the resonance hybrid is close to the more stable form than the less stable form.
more stable form(negative charge on
electronegative oxygen)
Ch.2 Polar Bonds and Their Consequences
4. Resonance forms must be valid Lewis structures and obey normal rules of valency; The octet rule still applies.
5. The resonance hybrid is more stable than any individual resonance form; The larger the number of resonance forms, the more stable a substance is.
C CO
O
H
H
H
Acetate ion (two resonance forms)
C CO
O
H
H
H
-C C
O
O
H
H
H
NOT a validresonance form
10 electrons on carbon
Ch.2 Polar Bonds and Their Consequences
X
Y
Z
*
XY
Z* *XY
Z X
Y
Z
*
* = 0, 1, or 2 electronsmultiple bond (double or triple bond)
X
Y
Z
2.6 Drawing Resonance Forms
Any three atom grouping with a multiple bond has two resonance forms.
OC
O
O_ _
Practice Resonance forms
OC
O
O_ _
OC
O
O_ _
OC
O
O_
_
OC
O
O_
_
Carbonate ion: CO32-
Ch.2 Polar Bonds and Their Consequences
Practice Resonance forms
CC
CC
CH
HH
H H
H
H
Pentadienyl radical
CC
CC
CH
HH
H H
H
HC
CC
CC
HH
H
H H
H
HC
CC
CC
HH
H
H H
H
H
allylic radical
A half-headed curved arrow shows movement of one electron.
Ch.2 Polar Bonds and Their Consequences
Ch.2 Polar Bonds and Their Consequences
2.7 Acids and Bases: The BrØnsted-Lowry Definition
BrØnsted-Lowry Definition: proton donor/acceptorLewis Definition: electron pair donor/acceptor
BrØnsted-Lowry Acid: proton (H+) donorBrØnsted-Lowry Base: proton (H+) acceptor
Ch.2 Polar Bonds and Their Consequences
+ H-B+ +
Acid Base Conjugateacid
Conjugatebase
H-A B A-
In general,
+ H2O +
Acid Base Conjugateacid
Conjugatebase
CH3COOH -OH CH3COO-
+ NH3 +
Acid Base Conjugateacid
Conjugatebase
H2O -NH2-OH
Ch.2 Polar Bonds and Their Consequences
2.8 Acid and Base Strength
+ H3O+ +H-A H2O A-
Keq =[H3O+][A-]
[HA][H2O]
Ka = Keq [H2O] =[H3O+][A-]
[HA]
Acidity constant, Ka
in dilute solution [H2O] remains constant [H2O] = 55.6 M for pure water
pKa = -logKa
stronger acid: larger Ka, smaller pKaweaker acid: smaller Ka, larger pKa
Ch.2 Polar Bonds and Their Consequences
Acid pKa Conjugate base
CH3CH2OHH2OHCNCH3COOHHFHNO3HCl
CH3CH2O-
HO-
CN-
CH3COO-
F-
NO3-
Cl-
16.0015.749.314.763.45- 1.3 - 7.0
strongeracid
strongerbase
Table 2.3 Relative strength of some common acids and their conjugate bases
Ch.2 Polar Bonds and Their Consequences
- inverse relationship between the acid strength and its conjugate base strength
- a strong acid loses H+ easily; its conjugate base has little affinity for H+
- a weak acid loses H+ with difficulty; its conjugate base has high affinity for H+
for example, HCl is a strong acid; Cl- does not hold H+ tightly and is thus a weak base
+ H-B+ +
Acid Base Conjugateacid
Conjugatebase
H-A B A-
Ch.2 Polar Bonds and Their Consequences
2.9 Predicting Acid-Base Reactions from pKa Values
An acid with a lower pKa will go to the conjugate acid with a higher pKa.
+ H2O+
Acetic acid(pKa = 4.76)
CH3COOH -OH CH3COO-
Water(pKa = 15.74)
Hydroxide ion Acetate ion
strongeracid
strongerbase
weakeracid
weakerbase
Practice Acidity
C C HH + HO- C CH + HO-H
pKa = 15.74pKa = 25
?
+
Acetone(pKa ~ 19)
H3C CH3
ONa+NH2
- ?
C C HH + NH2- C CH + NH3
pKa = 35pKa = 25
Ch.2 Polar Bonds and Their Consequences
+
Amonia(pKa ~ 35)
H3C CH2
ONH3
Na+
Ch.2 Polar Bonds and Their Consequences
2.10 Organic Acids and Organic Bases• organic acids are two kinds: O-H and O=C-C-H
H3C OH
H3C O
Anion is stabilized by havingnegative charge on a highly electronegative atom
H3C O
OH
H3C O
O
H3C O
O
Anion is stabilized by having negative charge on a highly electronegative atom and by resonance
pKa = 15.54
pKa = 4.76
Ch.2 Polar Bonds and Their Consequences
O H
H3C CH3
O O
H
pKa = 10 pKa = 9
H3C O
OH
pKa = 4.76
H3C CH2
OH
H3C CH2
O
H3C CH2
O
Anion is stabilized resonance and having negative charge on a highly electronegative atom
organic acids
pKa = 19.3
Ch.2 Polar Bonds and Their Consequences
Organic bases: only one main kind- nitrogen atom
H3C NHH
H-Cl + H3C NHH
HCl-
N
Pyridine
(CH3CH2)3N
Triethylamine
H3CO
H H3C CH3
O
Oxygen-containing compounds: act as bases with strong acidsalso act as acids (O-H, C-H) with bases
Ch.2 Polar Bonds and Their Consequences
2.11 Acids and Bases: The Lewis Definition
Lewis Acid: electron pair acceptorLewis Base: electron pair donor
B A+ B A
filledorbital
vacantorbital
Lewis base Lewis acid
Ch.2 Polar Bonds and Their Consequences
Cl H O HH
+ O HH
H+ Cl-
Lewis Acid / Base Reaction
BF
FF
O CH3CH3
+ BF
FF
OCH3
CH3
AlCl
ClCl
NCH3
CH3+ AlCl
ClCl
NCH3
CH3CH3
CH3
Ch.2 Polar Bonds and Their Consequences
Lewis acids:
neutral proton donors: H2O HCl HBr HNO3 H2SO4
CH3COOH CH3CH2OHOH
cations: Li+ Mg2+ Br+
metal compounds: AlCl3 BF3 TiCl4 FeCl3 ZnCl2 SnCl4
• Lewis definition of acid/base is broader than BrØnsted-Lowry definition: Lewis acid include many other species other than H+
Ch.2 Polar Bonds and Their Consequences
AlCl3 BF3tetrahedral complex
(maximum 4 CN)
Common Lewis Acids
OBBr3O
AlCl3
Ch.2 Polar Bonds and Their Consequences
Cl H O HH
+ O HH
H+ Cl-
Lewis base
Lewis Base: nonbonding electron pair
Ch.2 Polar Bonds and Their Consequences
Lewis bases:
CH3CH2OH CH3OCH3 CO
H3C H CO
H3C CH3
CO
H3C Cl CO
H3C OH CO
H3C OCH3 CO
H3C NH2
CH3SCH3NCH3
CH3H3C
• most oxygen and nitrogen containing compounds are Lewis bases
Ch.2 Polar Bonds and Their Consequences
CH3CH2OH CO
H3C OH
Some compounds, such as alcohols and carboxylic acids, can act as both acids and bases.
CH3CH2O CO
H3C O HHBrØnsted-Lowry Acid: proton (H+) donor
Lewis base: electron pair donor
Ch.2 Polar Bonds and Their Consequences
H3C CH3
O+ H2SO4
H3C CH3
O+ BF3
H3C CH3
O+
H
HSO4-
H3C CH3
OBF3
Ch.2 Polar Bonds and Their Consequences
H3C O
O+ H2SO4
H
- Reaction normally occurs only once in such instances, and the more stable of the two possible protonation products is formed.
• Some compounds, such as carboxylic acids, ester, and amides, have more than one atom with a lone pair of electrons and therefore react at more than one site.
For acetic acid, protonation occurs on the doubly bonded oxygen.
more stable(resonance stabilization) H3C O
O
H
H
+ HSO4-
H3C O
O
H
H
or H3C O
O
H
H
Ch.2 Polar Bonds and Their Consequences
CH
CH
HCH
H
HCH
HH
CH HH
CH3CH2CHCH3 CH3CH2CH(CH3)2
CH3or
condensed structure
2.12 Drawing Chemical Structures
condensed structure: C-H and C-C single bonds aren't shown
Ch.2 Polar Bonds and Their Consequences
H2CC
CH
CH2
CH3
H2CH2C
CH2
CH2
CH
H2C CH3 CH3
skeleton structure: line drawing of C-C bond- carbon atoms aren't usually shown- hydrogen atoms bonded to carbons aren't shown- atoms other than C, H are shown- C, H can be indicated for emphasis and clarity
Ch.2 Polar Bonds and Their Consequences
CC
CCC
CH
H
HOH
H
HOH
NHCH3HO
HO
OH
Adrenaline
HNHO
HO
OH
CH3
Space-filling model
Ball-and-stick modelWire-frame model
Molecular Models
Buckminsterfullerene (C60)