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Topic 4

Hydrides, Oxides, and Halides of thes- and p-block elements

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Hydrogen

• Hydrogen forms more compounds than any other element–> Three electronic possible processes:

(1) loss of a valence electron to give H+ (proton acids)(2) acquisition of an electron to give H– (hydrides)(3) formation of a covalent bond as in CH4

Note: There are many “in between cases”:– Formation of metallic hydrides (not regarded as simple ionic hydrides)– Formation of hydrogen bridge bonds in electron deficient compounds or

transition metal complexes

– Hydrogen bonding in polar solvents

BH

HB

H H

H HOC Cr H

CO

COOC

COCr

CO

CO

COOC

CO

2

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Hydrides

• All compounds of hydrogen could be termed “hydrides”, but not alldisplay “hydridic” character–> Hydridic substances react either as hydride ion (H–) donors orclearly contain anionic hydrogen:

• “Neutral” hydrogen compounds (e.g. CH4) have bonds with highlycovalent character

• “Acidic” hydrogen compounds have highly polarized bonds whichdissociate in polar solvents:

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Classification of Binary Hydrides

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Synthesis of Hydrogen Compounds

• Direct combination of elements:

• Protonation of a Brønsted base

• Metathesis (= double replacement) of a halide with a hydride

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Reactions of Hydrogen Compounds

• Heterolytic cleavage by hydride transfer:

• Heterolytic cleavage by proton transfer:

• Homolytic cleavage:

• Oxidation (gives EnOm and H2O, except metallic and group 17)

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4-7

Bond Energies

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Proton Affinity

• The proton affinity (Ap) is defined as the energy associated with theheterolytic cleavage of the E–H bond in the gas phase:

H+ + E– E–H

H• + E•

–Ap

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4-9

Proton Affinity Data

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Hydrides of Boron

• Diborane B2H6 is a gas (bp –92.6°C) that is flammable in air andinstantly hydrolyzed by H2O:

• Borane BH3 is unstable and is formed via thermal decomposition ofdiborane:

• The oxidation of diborane B2H6 is extremely exothermic:

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Structure and Bonding in Boranes

• The structure and bonding in boranes is unlike those of otherhydrides (e.g. CH4)

• There are not sufficient electrons to allow formation of conventionaltwo-electron bonds –> electron deficient bonds

• In the valence bond model the B–H–B bridge in diborane can bedescribed as a three-center-two-electron bond (3c-2e bond):

Ya

Yn

YbBoron sp3

hybrid orbitalsHydrogen 1s

orbitals

MOsenergy

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Neutral Boranes

Beside diborane there is a large number of borane cluster compounds:

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4-13

Structure and Bonding in Boranes

• To account for the structure and bonding in higher borane there area 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

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Valence Description of Boranes

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Hydrogen Storage

Holding hydrogen gas in your hands...Various metal alloys arecommercially available as asafe way to store hydrogen

Nanocrystalline hydrides show a very fast kinetics forthe absorption and desorption process:

Storage capacities for different methods:

H2 gas at 100 barLiquid H2

Hydride MgH2

0.81 mole/cm3

7.0 mole/cm3

11.1 mole/cm3

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Carbon Nanotubes

• Recently, carbon nanotubes have been showed to reversibly storehydrogen gas:

Graphite-like structureof carbon-nanotubes:

TEM image of a carbon-nanotube assembly:

Loading and unloading ofnanotubes with hydrogen:

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Classification of Binary Oxides

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Properties of Oxides

• Ionic oxides always react with water to give basic solutions

• The more covalent oxides that react with water always generateacidic solutions

• Polymeric oxides: there are basically two different types– Truly extended structures:– Molecular species such as P4O6 or P4O10 (and other group 15 analogs)

–> There is a great difference in reactivity depending on the nature oftheir “polymeric” structure

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Silicate Minerals

• Silicon forms a very large number of compounds containingtetrahedral SiO4

4– anions–> contained in many minerals (more than 80% of the earth’s crustatoms are silicon and oxygen!)

• Quartz is the most common form of silica (SiO2)

Crystal structure of b-Quartz

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Comm

on Silicate Structures

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Why is Silica Polymeric?

• Why is carbon dioxide monomeric (a molecular species), whereassilicon dioxide (=silica) is polymeric?–> the pi-bonds that involve p-orbitals other than 2p (3rd row andheavier elements) are weaker than two sigma bonds:

C–O 335 kJmol–1 x 2 = 670 kJmol–1

C=O 715 kJmol–1Si–O 420 kJmol–1 x 2 = 840 kJmol–1

Si=O 590 kJmol–1

• Although SiO2 does not react with water to generate an acidicsolution, its acidic properties are revealed by the reaction with base:SiO2 + 4 OH– SiO4

4– + 2 H2O

• SiO44– reacts with protons to generate condensed species:

SiO44–+ 2 H3O+ [O3Si-O-SiO3 ]6– + 3 H2O

• Further condensations lead to more highly polymerized species–> formation of silicate minerals

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Aluminosilicates

• Al3+ can replace Si4+ as long as electroneutrality is maintained bycompensating cations forming feldspars, zeolites and ultramarines.

–> both natural and synthetic zeolites find wide applications ascation exchangers (some cations will fit more snugly in the cavities),“molecular sieves” to absorb water or other uncharged smallmolecules (CO2, NH3, organic compounds)

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Zeolite Structures

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Classification of Oxides

Circles: amphoteric behaviorOctagons: amphoteric in lower oxidation states, acidic in higher

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Acid-base Properties of Oxides

• Basic oxides react with water to generate hydroxide:

• Similar reactivity is observed for group 1 and 2 salts of anions ofother p-block elements:

• Amphoteric oxides display both acidic and basic reactivity:

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Aluminum

• Aluminum is manufactured by electrolysis of alumina, which isobtained from Bauxite (1886, Hall-Heroult process)

– The ore is mixed with hot solution of NaOH, which will dissolve theoxides of aluminum and silicon but not other impurities

– The alumina is recovered by precipitation using CO2, an acidic oxide:

2 [Al(OH)4]– + 2 CO2 Al2O3 + 2 HCO3– + 3 H2O

– The purified alumina is mixed with cryolite (mixture of NaF and AlF3)and heated to 980°C to melt the solids

– The mixture is electrolyzed at 4-5 Volts and 50’000-150’000 Amps.–> one pound of aluminum requires 6-8 kilowatts of energy!

A 300 W light bulb can run for one hour with the energy it takes to makea single pop can...

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Electrolysis of Alumina

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Acidic Oxides

• Acidic oxides react with water to generate hydronium ions (H3O+):

SO2 + 2 H2O H3O+ + 2 HSO3–

• P4O10 is the acidic anhydride of phosphoric acid (H3PO4) and can beused to generate acid anydrides of other hydroxylic acids:

P4O10 + 12 HONO2 6 N2O5 + 4 H3PO4

–> N2O5 reacts with water to regenerate HONO2

(Note: NO and N2O are not acidic nitrogen oxides)

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Molecules Containing Sulfur and O

xygen

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Nitric Acid

• NO2 undergoes a redox process upon reaction with water:

• NO2 is synthesized from N2 and H2 as follows:

N2 + 3 H2 3 NH3

2 NH3 + 5/2 O2 2 NO + 3 H2ONO + 1/2 O2 NO2

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Molecules Containing Nitrogen and O

xygen

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Binary Halides• As for the hydrides and oxides, halides form ionic, polymeric and

molecular (covalent) binary compounds with other elements:

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Synthesis of Binary Halides

• Direct interaction with the elements:– Most elements can be directly reacted with halogens to form halides:

– For metals, HF, HCl, and HBr may also be used

• Treatment of oxides with other halogen compounds:– Oxides are often replaced by halogens-containing compounds:

Pr2O3 + 6 NH4Cl 3 PrCl3 + 3 H2O + 6 NH3

• Halogen exchange:

– Many halides react to exchange halogen with elemental halogens, acidhalides or halide salts–> this is especially important for the synthesis of fluorides fromchlorides

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Molecular Halides

• Most of the eletronegative elements, and metals in high oxidationstates form molecular halides–> gases, liquids or solids with molecules held together only bydispersion (van der Waals) forces

• Molecular halides are often easily hydrolyzed:BCl3 + 3 H2O B(OH)3 + 3 H+ + 3 Cl–

SiCl4 + 4 H2O Si(OH)4 + 4 H+ + 4 Cl–

F

P

F

FFFF B

F

F

FC

F

FF

F

S

F

FFF

F

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Bridged Halides

• Polymeric halide compounds are formed with elements ofintermediate electronegativity–> the partial positive charge on the bonding partner can furtherinteract with the partial negative charge on the halide formingbridged structures

SbFF

F

F

F

FSb

F

F

FF

FSb

F

F

F

F

F

TeFFF

FTe

FF

F

FTe

FF

F

FFS

FFF

F

F

P

F

FFF

ClBe

Cl

Cl ClBe

Cl

ClBe

ClCl

BeCl

Cl

ClGa

Cl

Cl ClGa

Cl

Cl

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Bonding in Bridged Halides• The bonding in bridged halides is similar as in boranes• The A–X–A bridge in bridged halides can be described as a

three-center-four-electron bond (3c-4e bond)• There are not sufficient bonding orbitals to allow formation of conventional two-

electron bonds –> orbital deficient bondNote: the bond order is only 0.5 (per A–X bond), since 2 electrons are in nonbondingorbitals!

Yb

MOs

Ya

Yn

sp3 hybridorbitals

halide porbital

energy

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Polyatomic Halide Ions

• In addition to the common monoatomic halide ions, numerouspolyatomic species (cationic and anionic) have been prepared:

I2 + I– I 3–

K = 698 at 25°C in aq. solution

Some examples ofpolyiodide ions

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Description of Bonding in I3–

• The linear structure of I3– can be described via a sp3d hybridizedcentral atom and a total of five electron pairs (two BPs, 3 LPs)

• Alternatively, the central atom can be looked at as sp2+p hybridizedresulting a three-center-four-electron bond–> the resulting bond order is 0.5, which readily accounts for theweaker axial bond found in molecules such as PX5, BrF3, etc.Note: I–I bond length in I3– 290 pm, in I2

267 pm!

energy

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Interhalogens

• Halogens form many compounds containing two or more differenthalogens–> these might be diatomic (ClF) or polyatomic (such as ClF3, BrF5or IF7)

Note the size effect of central atom: only iodine forms compounds with 7and 8 substituents

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Pseudohalogens

• Parallels have been observed between the chemistry of the halogensand a number of other dimeric species (= pseudohalogens):

• Neutral diatomic species• Ion of 1– charge• Formation of hydrohalic acids• Formation of interhalogen compounds• Insolubility in water of salts with heavy metals such as Ag+ and Pb2+

• Addition of halogens across multiple bonds

Cl2Cl–HClI!Cl, BrCl, ClFAgCl, PbCl2

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Noble Gas Compounds

• Before 1962 no compounds containing covalently bonded noblegases were known–> since then a large number of fluorides and oxides have beenprepared

• XeF2 (linear) and XeF4 (planar) have structures in accord with theVSEPR predictions, XeF6 and XeF8

2– are difficult to interpret

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Nobel Gas Fluorides

• Noble gas fluorides are formed by mixing xenon and fluorine andactivating the mixture by thermal or photochemical energy:

Xe + F2 XeF2

Xe + 2 F2 XeF4

Xe + 3 F2 XeF6

The challenge in thesereactions is the isolation ofpure compounds, since allthree products tend to form