CHEM 1311A Syllabus
• Some descriptive chemistry of s- and p-block elements
– Compounds of hydrogen
– Compounds of oxygen and related elements
– Compounds of the halogens (halides)
– Acids and bases
• Redox chemistry of the elements
Third Exam – Friday, October 31
Hydrogen Compounds
• Classification• Synthesis
– Direct combination of the elements– Protonation of a BrNnsted base– Metathesis (double replacement) of a halide with a
hydride• Reactions
– Homolytic cleavage– Heterolytic cleavage by hydride transfer– Heterolytic cleavage by proton transfer– Oxidation (gives EnOm and H2O except metallic and
group 17)
Classification of Hydrogen Compounds
Synthesis: direct combination of the elements
M(s) + H2(g)
X2(g) + H2(g)
salt-like metal hydrides are among the strongest bases known
N2(g) + H2(g)
Synthesis: protonation of a BrNnsted base
Mg3N2 + H2O
(3 Mg(s) + 2 N2(g) Mg3N2(s))heat
Ca3P2 + H2O
CaC2 + H2O
Other examples of Brϕnsted bases include Na2O, Na2S, CaH2, NaO2CR
Brϕnsted base – proton acceptor
Synthesis: metathesis (double replacement) of a halide with a hydride
LiH + AlCl3
NaH + BF3
LiAlH4 + SiCl4
LiAlH4 + BF3
BH
BH
HHHH
3-center-2-electron bonds
3-center-2-electron bonds
Many examples in boron and aluminum compounds; most often hydrogen is bridging but sometimes carbon groups, i.e., CH3 in compounds such as Al2(CH3)6 which is dimeric with bridging methyl groups
B2H7-, B2H6, B4H10, B5H9, B5H11, B6H10, B10H14
A MO description of the 3-center-2-electron bond in H3B-H-BH3- can
be generated assuming that each boron uses an sp3 hybrid orbital for combination with the hydrogen 1s orbital
E
Linear combinations of boron orbitals
Neutral Boranes
Besides diborane there is a large number of borane cluster compounds:
B4H10
B5H11
B5H9
B6H10
B10H14
Structure and Bonding in Boranes
• To account for the structure and bonding in higher borane there are a total of five structurally different bonding elements present:
Terminal boron-hydrogen bond
Hydrogen bridge bond
Boron-boron bond
Open B–B–B bond
Closed boron bond
2c–2e
3c–2e
2c–2e
3c–2e
3c–2e
B H
B
H
B
B B
B
B
BB
BB
Bonding element Bonding type Symbol
Definitions of acids and bases
● Arhennius: H+ is the acid and OH- is the base.
● Bronsted: An acid is a proton donor and a base is a proton acceptor.
● Lewis: An acid is an electron pair acceptor and a base is an electron pair donor.
Solvent leveling effect on acid/base strengths
● The strongest acid and the strongest base that can exist in a given solvent are the acid/base (conjugate acid/conjugate base) generated by autoionization of the solvent.
2 H2O = H3O+ + OH-
2 NH3 = NH4+ + NH2
-
2 CH3CO2H = CH3C(OH)3+ + CH3CO3
-
● Intrinsically stronger acids or bases will be “leveled” to the strength of the conjugate acid or conjugate base of the solvent.
NaNH2 + H2O(liq) = NH3(sol) + Na+(sol) + OH-
(sol)
HCl + NH3(liq) = NH4(sol)+ + Cl-(sol)
Reactions: homolytic cleavage
E-H EC + HC
Thermal cleavage will occur at moderate temperatures only for very weak element (E) to hydrogen bonds.
I-H295
Te-H268
Sb-H257
Sn-H264
In-H243
Br-H362
Se-H276
As-H247
Ge-H288
Ga-H<274
Cl-H428
S-H363
P-H322
Si-H318
Al-H285
F-H565
O-H459
N-H386
C-H411
B-H389
E-H bond energies (average), kJ mol-1
Heterolytic cleavage by hydride transfer
RXH is a Brϕnsted acid (X is an electronegative element)
NaH + H2O
MH + HXR
E-H + H+
SiH4 + 2 H2O
Heterolytic cleavage by proton transfer
H2O W
HX + H2O W
CH3CH2OH + NaH
HX + OH-
PH3 + NaH
PH3 + NaNH2
Oxidation of hydrogen compounds gives EnOm
and H2O (except metallic and group 17)
CH4 + 2 O2 = CO2 + 2 H2O
SiH4 + 2 O2 = SiO2 + 2 H2O
4 PH3 + 6 O2 = P4O6 + 6 H2O and P4O6 + 2 O2 = P4O10
B4H10 + 11/2 O2 = 2 B2O3 + 5 H2O
Dilute solutions of alkali metals in liquid NH3
• Dilute solutions are deep blue with an absorption spectrum that is independent of the metal.
• Show conductivities that are comparable to those of electrolytesof equivalent concentration.
• Have magnetic susceptibilities consistent with ca. one unpaired electron per sodium.
• Exhibit electron spin resonance spectra that are consistent withthose of free electrons and inconsistent with metal-based electrons
– M(s) W M(sol)+ + e(sol)
-
– 2 M(s) W M(sol)+ + M(sol)
-
– M(sol)- W M(sol) + e(sol)
-
Conc. solutions of alkali metals in liquid NH3
• Bronze colors with metallic luster.• Metallic conductivities.• Magnetic behavior comparable to that of pure metals.• M(sol)
Solutions of alkali metals in other solvents
• Solubility varies with identity of solvent, but generally only dilute solutions can be made in absence of chelating agent.
• Metal independent spectral features are weak, when present, and are invariably dominated by metal dependent features. Only the latter are observed in the presence of chelating agents
• Sodium is insoluble in some solvents, but K-Na alloy dissolves.– K(s) + Na(s) W K(sol)
+ + Na(sol)-
Isolation of salts of the sodium anion
• This occurs because of the equilibrium 2 Na(S) + C2.2.2 W[Na(C2.2.2)]+ + Na&
• Crystallization at low temperature gives [Na(C]2.2.2)]Na whose crystal structure is identical to that of [NaC]I except that the radius of Na& appears to be somewhat greater than that of iodide.
• 23Na NMR measurements on [Na(C]2.2.2)]Na in solution indicate two types of sodium, one identical to that in [Na(C]2.2.2)]X andanother consistent with predictions for Na&.
OO
OO
NN
OO
• Sodium is only slightly soluble in ethylamine (10-6 M) but gives 0.2 M solutions in the presence of C2.2.2.
C2.2.2
Definition of proton affinity
HC = H+ + e- 1312 kJ mol-1
NH2C + e- = NH2- -75 kJ mol-1
NH3 = NH2C + HC 454 kJ mol-1
NH3 = NH2- + H+ 1691 kJ mol-1
Ap = IE - EA + Bdiss
Proton affinity, Ap, is defined as the energy associated with the heterolytic cleavage of the E-H bond in the gas phase
H+ + E- E-H
+H E
-Ap
-IE EA
-Bdiss••
Proton affinities for EHn-*
Proton affinity is for reaction H-EHn = H+ + EHn-
I-
1315
Br-
1354
SeH-
1466
AsH2-
1502
GeH3-
1509
Cl-
1395
SH-
1476
PH2-
1552
SiH3-
1554
F-
1554
OH-
1635
NH2-
1689
CH3-
1745
-200
-100
0
100
200
300
400
H Li Be B C N O F Na Mg Al Si P S Cl K Ca
EA
/kJ m
ol-1
E-H bond energies (average), kJ mol-1
I-H295
Te-H268
Sb-H257
Sn-H264
In-H243
Br-H362
Se-H276
As-H247
Ge-H288
Ga-H<274
Cl-H428
S-H363
P-H322
Si-H318
Al-H285
F-H565
O-H459
N-H386
C-H411
B-H389
CH4 = CH3 + H 435CH3 = CH2 + H 443CH2 = CH + H 443CH = C + H 339Average 415
Step-wise dissociation energies for methane
Inductive effects on acid strengths
pKa
Cl-OH 7.5Br-OH 8.7I-OH 10.7HO-OH 11.8H-OH 15.7CH3-OH 16.6
F-CH2C(O)OH 2.7 F-SO2-OH > Cl-SO2-OH > HO-SO2-OHCl-CH2C(O)OH 2.8Br-CH2C(O)OH 2.9 F2P(O)(OH) > FP(O)(OH)2 > HOP(O)(OH)2
I-CH2C(O)OH 3.0H-CH2C(O)OH 4.7CH3-CH2C(O)OH 4.9
Similar effects exist for bases. For example base strengths vary in the order NH3 > H2NNH2 > ClNH2 > Cl2NH > F3N, etc.
pKa values for hydrated metal ions, [M(OH2)n]m+
Mm+ pKa Mm+ pKa
Th4+ 3.2
Al3+ 5.0 Sc3+ 4.3
Y3+ 7.7 Cr3+ 4.0
La3+ 8.5 Fe3+ 2.2
Mg2+ 11.4 Cr2+ 10.0
Ca2+ 12.8 Mn2+ 10.6
Sr2+ 13.3 Fe2+ 9.5
Ba2+ 13.5 Co2+ 9.6
Ni2+ 9.9
Zn2+ 9.0
Li+ 13.6 Ag+ 12.0
Na+ 14.2 Tl+ 13.2
K+ 14.5
pKa=-log K for [M(OH2)n]m+ + H2O = [M(OH2)n-1OH](m-1)+ + H3O+
smaller pKa = greater dissociation, stronger acid
pKa values for acids of type (HO)mEOn
(oxidation state effects)
10.7HOI
8.7HOBr
3
2
1
0
No. E=O
11.66.772.25(HO)3AsO
-10HOClO3
1.92-3(HO)2SO2
-1.4HONO2
-1HOClO2
12.37.212.16(HO)3PO
3.3HONO
1.94HOClO
7.4HOCl
pK3pK2pK1Acid
Solvent leveling (of acidities and basicities)
• The strongest acid and strongest base that can exist in a given solvent are the conjugate acid and base generated by autoionization
– 2 H2O = H3O+ + OH-
– 2 NH3 = NH4+ + NH2
-
– 2 CH3CO2H = CH3C(OH)2+ + CH3CO2
-
– 2 HF = H2F+ + F-
– 2 BrF3 = BrF2+ + BrF4
-
• An intrinsically stronger acid/base will be leveled (reduced) to the acidity/basicity of the conjugate acid or base for the solvent.
• To determine intrinsic acid/base strengths a solvent must be used that is sufficiently low in basicity/acidity such that the acid or base under investigation is not completely ionized.
• Solvation of ions affects ionization constant.
Acidities measured in CH3CNHo = pKBH+ - log [BH+]/[B]H2SO4 -2HCl -7HBr -9HClO4 -10HI -11
NH2
NO2
NO2
pKBH+ = -4.5
Superacids (no solvent)
(Ho = pKBH+ - log [BH+]/[B])HF -11H2SO4 -11.9HClO4 -13CF3SO3H -14.6FSO3H -15.6FSO3H/SbF5 (25%) -21 (Magic Acid)HF/SbF5 (1:1) -28 (est.)Olah awarded Nobel Prize in Chemistry in 1994
Applications of superacids
FSO3H/SbF5or HF/SbF5
R2OROHRX
RH
RC
H=C
H2
ArH
R2CO
RCHO
RSH
R2S
RCO2H
RC
(O)O
R
RC
(O)N
R2
(RO) 2
CO
R2OH+ROH2+
R+ + HX
R+ + H2
RCH+CH3
ArH2+
R2COH+
RCHOH+
(RO)2COH+
RSH2+
R2SH+
RC(OH)2+ RCO+ + H2O
RC+(OH)ORRC+(OH)NR2
Oxides: structural classification
PoO2Bi2O3
Bi2O5
PbO
PbO2
Tl2O
Tl2O3
BaOCs2O
Molecular CovalentPolymericIonic
XeO3
XeO4
I2O4
I2O5
I2O9
TeO2
TeO3
Sb4O6
Sb2O5
SnO
SnO2
In2O3SrORb2O
Br2O
BrO2
SeO2
SeO3
As4O6
As2O5
GeO2Ga2O3CaOK2O
KO2
K2O2
Cl2O
ClO2
Cl2O7
SO2
SO3
P4O6
P4O10
SiO2Al2O3MgONa2O
NaO2
Na2O2
F2O
F2O2
O2
O3
NO
N2O
N2O3N2O4
N2O5
CO
CO2
B2O3BeOLi2O
LiO2
Li2O2
Ionic vs covalent character and electronegativity
van Arkel-Ketalaar triange, p. 57 Norman
Acid-base properties of s- and p-block oxides
acidicbasiccircle – amphotericoctagon - amphoteric in lower oxidation state, acidic in higher
Sb
Po
Te
Se
S
BiPbTlBaCs
ISnInSrRb
BrAsGeGaCaK
ClPSiAlMgNa
NCBBeLi
H
Ionic oxides: bases
Na2O + H2O =
CaO + H3O+ =
Al2O3 + H3O+ + H2O =
Amphoteric oxides
Al2O3 + OH- + H2O =
Al(OH)4- + CO2 =
Acidic oxides
Group 13
B2O3 + 3 H2O = 2 B(OH)3
B(OH)3 + 2 H2O = H3O+ + B(OH)4-
Group 14
CO2 + H2O = (HO)2CO
(HO)2CO + H2O = H3O+ + HOCO2-
Overall: CO2 + 2 H2O = H3O+ + HOCO2-
SiO2 + 4 OH- = SiO44- + 2 H2O
SiO44- + 2 H3O+ = O3Si-O-SiO3
6- + 3 H2O
CO2 + OH- = HOCO2-
HOCO2- + OH- = CO3
2-
Overall: CO2 + 2 OH- = CO32- + H2O
LUMO for CO2
Silicates
SiO2 (cristobalite)
disilicate, Si2O76-
tetrahedral unit
infinite chain, SiO32-
pyroxenes
infinite double chain, Si4O11
6-, amphiboles infinite sheet, Si2O5
2-
micas
Si3O96-
Tremolite asbestos from Jamestown, CA
http://www.epa.gov/swerrims/ahec/summary/presentations/day1/addison1.pdf
Acidic oxides, con’t
Group 15
NO2 + NO = N2O3
2 NO2 = N2O4
2 NO2 + 2 OH- = NO3- + NO2
- + H2O
N2 + 3 H2 = 2 NH3
2 NH3 + 5/2 O2 = 2 NO + 3 H2O
NO + 1/2 O2 = NO2
2 NO2 + H2O = HONO2 + HONO
3 HONO = HONO2 + 2 NO + H2O recycle the NO
N2O5 + H2O = 2 HONO2
P4O6 + 6 H2O = 4 H3PO3
P4O10 + 6 H2O = 4 H3PO4
H3PO4 + H2O = H3O+ + H2PO4-
PP
P
PP O
P O P
OO
P OO
P O
P O P
OO
P OOO
O
O
O
Acidic oxides, con’t
Group 16
SO2 + H2O = H2SO3
H2SO3 + H2O = H3O+ + HSO3-
Overall: SO2 + 2 H2O = H3O+ + HSO3-
SO2 + OH- = HOSO2-
HOSO2- + OH- = SO3
2-
Overall: SO2 + 2 OH- = SO32 - + H2O
SO3 + H2O = H2SO4
H2SO4 + H2O = H3O+ + HSO4 -
HSO4 - + H2O = H3O+ + SO4
2-
Overall: SO3 + 3 H2O = 2 H3O+ + SO42-
solid SO3
SO
O
OO
O
SO
S
O OO
Group 17
Cl2O7 + H2O = 2 HOClO3
HOClO3 + H2O = H3O+ + ClO4-
I2O5 + H2O = 2 HOIO2
HOIO2 + H2O = H3O+ + IO3-
Structural classification of fluorides
171615141321
BiF3
BiF5
PbF2
PbF4
TlF
TlF3
BaF2CsF
IF
IF3
IF5
IF7
TeF4
TeF6
SbF3
SbF5
SnF2
SnF4
InF
InF3
SrF2RbF
BrF
BrF3
BrF5
SeF4
SeF6
AsF3
AsF5
GeF2
GeF4
GaF3CaF2KF
CIF
CIF3
CIF5
SF2
SF4
SF6
PF3
PF5
SiF4AlF3MgF2NaF
F2OF2NF3CF4BF3BeF2LiF
ionic, polymeric, molecular
Structural classification of chlorides
171615141321
BiCl3PbCl2TlCl
TlCl3
BaCl2CsCl
ICl
ICl3ICl5
TeCl4SbCl3SbCl5
SnCl2SnCl4
InCl
InCl3
SrCl2RbCl
BrCl
BrCl3
SeCl4AsCl3AsCl5
GeCl2GeCl4
GaCl3CaCl2KCl
CI2SCl2SCl4
PCl3PCl5
SiCl4AlCl3MgCl2NaCl
FClOCl2NCl3CCl4BCl3BeCl2LiCl
ionic, polymeric, molecular
Bonding in Bridged Halides
• The bonding in bridged halides appears similar to that in boranes; however,
• In group 13 compounds there are actually plenty of electrons andorbitals and the bonding is not electron deficient
2 BF3 + F-
same as
BF3 + BF4-
Description of bonding in I3–
• When there is an expansion of valence shells the situation is different
• The linear structure of I3– can be described via a sp3d hybridized central atom and a total of five electron pairs (two BPs, 3 LPs)
• Alternatively, the central atom can be looked at as sp2+p hybridized withonly the p orbital used to bond to the terminal iodine atoms– the resulting bond order is 0.5, which readily accounts for the weaker
axial bond found in molecules such as PX5, BrF3, etc.Note: I–I bond length in I3– 290 pm, in I2 267 pm!
energy
Description of bonding in I3– , cont’d
• Note that the situation is not improved by invoking a model that employs the s orbital on the central atom.
2 I• + I-
same as
I2 + I-
antibonding
bonding
bonding
Common reactions of covalent halides
As Lewis acid EXn + :B = B-EXn
B = electron pair donor; may be neutral or anionic
Common reactions of covalent halides
E-X + X- = X-E-X- E-X + H-OH = E-OH + H-X
E-X + H-OR = E-OR + H-XE-X + A = E+ + A-X-
Hard and soft acids and bases• First recognized for halide bases; acid behavior was referred to as class
a and class b– Kf for adduct with hard acids (class a) increases in order I- < Br- < CI-
< F-
• Hard acids include Al(III), Sc(III), Cu(II), Zn(II)– Kf for adduct with soft acids (class b) increases in order F- < Cl- < Br-
< I-
• Soft acids include Ag(I), Cd(II), Hg(II), Pb(II), Pd(II)• Later extended to many other bases
– Hard acids bond in order: R3P << R3N, R2S << R2O– Soft acids bond in order: R3N << R3P, R2O << R2S
• It follows that hard acids prefer hard bases and soft acids prefer soft bases– Hard-hard interactions are substantially electrostatic in nature; hard
acids are generally species with high energy LUMO’s and hard bases generally have low energy HOMO’s
– Soft-soft interactions are substantially more covalent in nature; soft acids and bases are generally larger and are significantly more polarizable. Soft acids generally have low energy LUMO’s and soft bases have high energy HOMO’s
More examples of hard and soft acids and bases
Bases
Acids
H2S, R2S, I-, SCN-, R3P, CN-, CO, H-, R-
NO2-, SO3
2-, Br-, N3
-, N2, C6H5N, SCN-
F-, OH-, H2O, NH3, CO3
2-, NO3-, O2-,
SO42-, PO4
3-, ClO4
-
Tl+, Ag+, BH3, Hg+, Hg2+, Ga(CH3)3, I2
Fe2+, Co2+, Ni2+, Cu2+, Zn2+, Pb2+, SO2, BBr3
H+, Li+, Na+, Be2+, Mg2+, Cr2+, Cr3+, Al3+, SO3, BF3, Al(CH3)3
SoftBorderlineHard
Spectral changes when I2 reacts with bases
E
s s
p p
I2
pB
I2