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