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 – Wednesday, March 25

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Today is also the deadline for dropping individual courses (W “grade” on transcript)

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

Magnesium Nitride Serves Up AmmoniaMagnesium nitride (Mg3N2) is a convenient source of ammonia for a host of reactions that produce nitrogen-containing compounds, a University of Cambridge research team reports (Org. Lett., DOI: 10.1021/ol801398z). The stable low-cost commercially available reagent, which produces NH3 when treated with water or certain alcohols, is an alternative to premade NH3 solutions or gas cylinders that can be cumbersome to use on a small scale. Gemma E. Veitch, Katy L. Bridgwood, and Steven V. Ley demonstrated the utility of Mg3N2in methanol as a reagent for the direct conversion of a broad range of esters to primary amides (carboxamides). In a separate study, the researchers used Mg3N2 to prepare dihydropyridines via the Hantzsch condensation of ethyl acetoacetate with various aldehydes in ethanol (Org. Lett., DOI: 10.1021/ol801399w). Magnesium nitride was previously known to release ammonia in water, the researchers note, but it had not been used to generate ammonia in organic syntheses until now. The team is still exploring the scope of the method, Bridgwood says. A paper on the preparation of pyrroles will be published soon, she adds, and experiments involving pyrimidine synthesis from nitriles, nucleophilic aromatic substitution, and continuous-flow reactions are ongoing.

Chem. Eng. News 2008 86, 36-37

Veitch, G.E.; Bridgwood K.L.; Ley, S.V.Org. Lett. 2008 10, 3623

Use of Mg3N2 with proton donor in synthesis

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 bondsMany 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

B B

Bonding element Bonding type Symbol

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 Na+ + OH- + H2

MH + HXR M+ + RX- + H2

E-H + H+ E+ + H2

SiH4 + 2 H2O SiO2 + 4 H2

Heterolytic cleavage by proton transfer

H2O W

HX + H2O W

CH3CH2OH + NaH

HX + OH-

PH3 + NaH

PH3 + NaNH2

Heterolytic cleavage by proton transfer

2 H2O W H3O+ + OH-

HX + H2O H3O+ + X-W

CH3CH2OH + NaH H2 + CH3CH2O- + Na+

HX + OH- H2O + X-

PH3 + NaH H2 + PH2- + Na+

PH3 + NaNH2 NH3 + PH2- + Na+

Oxidation of hydrogen compounds gives EnOmand 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

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.

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-

1354SeH-

1466AsH2

-

1502GeH3

-

1509

Cl-

1395SH-

1476PH2

-

1552SiH3

-

1554

F-

1554OH-

1635NH2

-

1689CH3

-

1745

AP vs pKa; intrinsic basicity vs effective basicity

AP BH+ = B + H+

Ka BH+ + H2O = B + H3O+

pKa = - log Ka

(smaller number, greater dissociation, weaker base)

Base AP, kJ mol-1 pKa

NH3 866 9.24

CH3NH2 904 10.6

(CH3)2NH 929 10.8

(CH3)3N 950 9.8

Inductive effects on acid strengthspKa

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)2I-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+ 3.2Al3+ 5.0 Sc3+ 4.3Y3+ 7.7 Cr3+ 4.0La3+ 8.5 Fe3+ 2.2Mg2+ 11.4 Cr2+ 10.0Ca2+ 12.8 Mn2+ 10.6Sr2+ 13.3 Fe2+ 9.5Ba2+ 13.5 Co2+ 9.6

Ni2+ 9.9Zn2+ 9.0

Li+ 13.6 Ag+ 12.0Na+ 14.2 Tl+ 13.2K+ 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)

10.7HOI8.7HOBr

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)3PO3.3HONO

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 conjugatge 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.

− NaNH2 + H2O(liq) = NH3(sol) + Na+(sol) + OH-

(sol)

− HCl + NH3(liq) = NH4(sol)+ + Cl-(sol)

Solvent leveling (of acidities and basicities)

● 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

R2OROHRXRH

RCH

=CH

2

ArH

R2CO

RCHO

RSH

R2S

RCO2HR

C(O)O

R

RC(

O)N

R 2

(RO) 2CO

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

PbOPbO2

Tl2OTl2O3

BaOCs2O

Molecular CovalentPolymericIonic

XeO3

XeO4

I2O4

I2O5

I2O9

TeO2

TeO3

Sb4O6

Sb2O5

SnOSnO2

In2O3SrORb2O

Br2OBrO2

SeO2

SeO3

As4O6

As2O5

GeO2Ga2O3CaOK2OKO2

K2O2

Cl2OClO2

Cl2O7

SO2

SO3

P4O6

P4O10

SiO2Al2O3MgONa2ONaO2

Na2O2

F2OF2O2

O2

O3

NON2ON2O3N2O4

N2O5

COCO2

B2O3BeOLi2OLiO2

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

S OO

OO

O

S O

SO O

O

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

TlFTlF3

BaF2CsF

IFIF3

IF5

IF7

TeF4

TeF6

SbF3

SbF5

SnF2

SnF4

InFInF3

SrF2RbF

BrFBrF3

BrF5

SeF4

SeF6

AsF3

AsF5

GeF2

GeF4

GaF3CaF2KF

CIFCIF3

CIF5

SF2

SF4

SF6

PF3

PF5

SiF4AlF3MgF2NaF

F2OF2NF3CF4BF3BeF2LiF

ionic, polymeric, molecular

Structural classification of chlorides

171615141321

BiCl3PbCl2TlClTlCl3

BaCl2CsCl

IClICl3ICl5

TeCl4SbCl3SbCl5

SnCl2SnCl4

InClInCl3

SrCl2RbCl

BrClBrCl3

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 and

orbitals and the bonding is not electron deficient – can be viewed as a 3-center-4-electron bond

2 BF3 + F-

same as

BF3 + BF4-

Figure adapted from Munzarova, M.L.; Hoffmann, RJ. Am. Chem. Soc. 2002, 124, 4784

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 with

bonding only the p orbital used to form bonds 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

Figure adapted from Munzarova, M.L.; Hoffmann, RJ. Am. Chem. Soc. 2002, 124, 4784

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-

How to recognize a Lewis acid or Lewis base

● Molecules for which a simple Lewis structure indicates an atom does not have four pairs of electrons in its valence shell (an octet) will behave as Lewis acids.

● Molecules whose Lewis structures indicate an atom to have an octet as a result of the formation of one or more multiple bonds will often function as Lewis acids.

● Molecules that have central atoms that are capable of expanding their valence shell to include more than four pairs of sigma-bonding electron pairs, i.e., use hybrid orbitals that involve d orbitals, are Lewis acids.

● The acceptor orbital is always the LUMO.

● Lewis bases are molecules with lone pairs of electrons.– In some molecules the energies of the orbitals containing lone pairs are

too low in energy to interact appreciably with most Lewis acids.– In some instances bonding F or B electrons are basic.– Often the HOMO is the donor orbital but the nature of the acid can result

in a lower energy orbital functioning as the donor orbital.

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