course on inorganic chemistry by frank klose chapter 1

Course on Inorganic Chemistry by Frank Klose

Chapter 1

Elements and Compounds, Atoms and Molecules – Structures and Bonds

Substances, Compounds and Elements

Substances

Homogeneous substances

Heterogeneous substances

Solutions Pure substances

Compounds Elements

Substances

Homogeneous substances

Heterogeneous substances

Solutions Pure substances

Compounds Elements

The Discovery of the Chemical Elements

Antiquity/Middle Ages - The “Four Elements” Fire, Water, Earth, and Air

1642 – Jungius 1661 – R. Boyles 1777 – Lavoisier

- Pioneer works on the present-day theory of chemical elements, definition of the terms “element” and “compound” (Lavoisier)

1869 - Proposal of the “Periodic Table of Elements” by Mendelejew and Meyer (independently)

1999 - Discovery of the elements 114 (Joint Institute for Nuclear Research Dubna, Russia), 116, 118 (Berkeley, California, USA)

Number of elements discovered

12 14 1532

82

115

175

0

50

100

150

200

Antiquity

13th century

17th century

18th century

19th century

20th century

Theoretical maximum

AntiquityAntiquity

13th century13th century





Theoretical maximumTheoretical maximum

Percentage of Elements

Earth’s crust Human “biomass” Element Percentage

[mg/kg]

Percentage

Element [mg/kg]

O 467600 O* 611000 Si 278600 C* 236000 Al 81600 H* 94000 Fe 50200 N* 28000 Ca 36400 Ca* 14000 Na 28400 P* 9330 K 26000 S* 2330

Mg 21000 K* 2270 Cl 18100 Na* 1400 Ti 4400 Cl* 1400 H 1400 Mg* 440 P 1000 Fe* 56

Mn 950 Zn* 40 F 625 Si* 18.7

Ba 425 Rb 18.7 Sr 375 F* 10.7 S 260 Sr 4 C 200 Zr 4 N 20 Cu* 2.67 Zr 165 Br 1.87 V 135 Sn* 1.87 Cr 100 Nb 1.33

I* 0.933 Al 0.467 Pb 0.467 Cd 0.4

Rb, Ni, Zn, Ce, Cu, Y, La, Nd, Co, Sc, Li, Nb,

Ga, Pb, Th, B

10 - 100

Ba 0.267 Mn* 0.267

1 - 10 V* 0.267 B 0.187 Se* 0.187 Mo* 0.0667

Pr, Br, Sm, Gd, Ar, Yb, Cs, Dy, Hf, Er, Be, Xe, Ta, Sn, U, As, W, Mo,

Ge, Ho, Eu

As* 0.0467 0.1 – 1 Co* 0.0373

Cr* 0.0267 Tb, I, Tl, Tm, Lu, Sb,

Cd, Bi, In Li 0.0267

< 0.1 Ni* 0.0133

Hg, Ag, Se, Ru, Te, Pd, Pt, Rh, Os, Ne, He, Au, Re, Ir, Kr, Ra, Pa, Ac,

Po, Rn, Np, Pu, Pm, Fr, At, Transplutonium

elements

*) essentially

The Atom and its Components

1808 - Hypothesis of atoms by Dalton 1897 - Discovery of the electron by J. J. Thomson 1913 - Discovery of the proton by E. Rutherford 1911 - Electron scattering experiments by

E. Rutherford – Atom model concept proposing a dense positive charged core and a negative charged but near mass less shell filled with electrons

1921 - Discovery of the neutron by W. D. Harkins 1926 - Schrödinger equation, begin of the quantum

mechanical description of atoms 1932 - Atom model by W. Heisenberg using electron

orbitals

Atom

ShellCore

Protons Neutrons

negative charged

Electrons

positive charged

nocharge

contain the mass of an atom

responsible for chemical properties (outer e-)

Atom

ShellCore

Protons Neutrons

negative charged

Electrons

positive charged

nocharge

contain the mass of an atom

responsible for chemical properties (outer e-)

Core/shell ratios :

- 10-4 with respect to the radius - 5000 : 1 with respect to the mass

(99.95 – 99.98 % of the atom mass is concentrated in the core)

Atomic Constants and Dimensions

Masses absolute mass of a proton: 1.6726 * 10-27 kg absolute mass of a neutron: 1.6749 * 10-27 kg absolute mass of an electron: 9.1093 * 10-31 kg absolute mass unit u [amu] = 1/12 * m(12C) = 1.6605 * 10-27 kg

Relative masses of atoms (Ar) and molecules (Mr)

Ar or Mr = (mA or mM)/u (IUPAC 1961)

Molar masses M [g/mol] M = m * NA

→ Numbers of Ar or Mr and M are identically!

Radius of atoms: 0.3…3 * 10-10 m (10-10 m = 1 Å (Angstroem)) Other important constants e - elementary charge: 1.6022 * 10-19 C

NA - Avogadro number: 6.0221 * 1023

h - Planck constant: 6.6261 * 10-34 J * K-1

Fundamental Equations from Quantum Mechanics

Schrödinger equation (1926)

H ψ= E ψ

H – Hamilton operator E – Energy of the electron

ψ - Wave function The Uncertainty Principle by Heisenberg (1927)

∆x * ∆p ≥ h/4π

∆x – uncertainty of the position of the electron ∆p - uncertainty of the impulse of the electron

h – Planck constant

Electron orbitals as the solutions of the Schrödinger equation:

→ rooms of highest probability (90 % or more) of finding the electron → motion of electrons in the orbitals is free of energy loss → electron energy levels are discrete

Electron Orbitals of Atoms

s orbital

px orbital

py orbital

pz orbital

dxy orbital dxz orbital dvz orbital

22 yx

d− orbital

2zd orbital

Algebraic signs are related to the angular part of the wave function, not to a charge!

Quantum Numbers for Electron Orbitals

The three fundamental properties of electrons: mass, charge, spin The Pauli principle (Wolfgang Pauli, 1924):

No more than two electrons can occupy any given orbital. If two electrons do occupy one orbital, then their spins must be paired. Every electron orbital can be characterised by a set of quantum numbers definitely.

n – principal number - determines the number of the shell (n = 1, 2, 3, …) - sometimes shells named with capital letters K, L, M, …

(e.g. X-ray analysis) l – orbital angular momentum quantum number (subshell number)

- determines the type of electron orbital (s, p, d, f, g, …) - l can range from 0 to (n - 1) - number of orbitals of a shell n is n²

m – orbital magnetic quantum number

- determines the orientation of the orbital (x, y, z, …) - unoccupied orbitals differing in m have the same energy

(they are “degenerated”) - energy split in many electron systems (coupling of angular and

magnetic momentum, Coulomb interactions) - m = 0, ±1, ±2,…, ±l,

s – spin magnetic quantum number

- the only values: - ½ , + ½ additionally: j – angular momentum quantum number

- j = l ± s, (all possible combinations of l and s)

The Energy Scheme for Electron Orbitals

n = 7

7 p

7 d

7 s

n = 7

7 p

7 d

7 s

Building up principals:

- Electrons occupy shells and orbitals in order of their energies (defined by n and l).

- Each inner shell should be fully filled before occupying the next shell. - Fully occupied subshells (s2, p6, d10, f14) have the highest stability. Half

occupied d subshells (d5) are favoured, too. - Hund’s rule: An atom in its ground state adopts a configuration with

the greatest number of unpaired electrons. Electrons occupy different orbitals of a given subshell before doubly occupying any one of them (Σs is maximised).

- Outer electron configuration (valence electrons) determines chemical properties.

- RESULT: Periodicity of number of valence electrons by sequential filling of s, p, d and f orbitals

Energy

Periodicity - The Size of Atom Orbitals

electron

core

1st shell2nd shell

3rd shell

electroncore

1st shell2nd shell

electron

core

1st shell2nd shell

3rd shell

electroncore

1st shell2nd shell

Radius of orbitals of neutral atoms Radius of orbitals of neutral atoms

→ contraction with increased proton number for each shell (increased Coloumb attraction between the positive charged core and the negative charged electron shell)

→ Positive ions are smaller and negative ions are larger compared to the neutral atom.

→ Energy of orbitals is specific for each element.

Periodicity – Ionisation Energies

First and second ionisation potential

Electron affinity

Atomic Spectroscopy

Principle of spectroscopy

Excitation (specific or non-specific),

absorption

→

Relaxation, Emission

of specific radiation Atom Absorption Spectroscopy (AAS)/ Optical Emission Spectroscopy (OES, OES-ICP) → using of outer electron transitions

(∆l = ±1, ∆j = 0, ±1, s → p and p → d transitions)

→ specific for each element → laser technology

Term scheme for sodium (Na)

X-ray Flourescence Spectrometry (XFS)

→ using of inner electron transitions (∆l = ±1, ∆j = 0, ±1, s → p and p → d transitions) → XFS: primary relaxation, applicable for elements with Z = 9 - 92

primary X-ray radiatation

secondary X-ray radiatation

electron energy [eV]

primary X-ray radiatation

secondary X-ray radiatation

electron energy [eV]

(1) absorption of primary X-ray radiatation → remove of a inner electron → ionisation

(2) transfer of an electron from an outer shell to the leak

(3) emission of secondary X-ray radiation (specific for the element)

The Periodic Table of Elements (PTE)

Ia IIa IIIa IVa Va VIa VIIa VIIIa

IIIb IVb Vb VIb VIIb VIIIb Ib IIb

Main group elements

d block elements(transition metals)

Lanthanides

Actinides

s block(l = 0)

d block(l = 2)

p block(l = 1)

n

Ia – VIIIa, Ib – VIIIb = group numbers 1-7 = numbers of periods

name of the element

chemical symbol

rel. atomic mass(= molecular mass)

atomic number Z (= number of protons)

electron negativity

colour = metallic or non-metallic character oracid-base properties of the oxides

name of the element

chemical symbol

rel. atomic mass(= molecular mass)

atomic number Z (= number of protons)

electron negativity

colour = metallic or non-metallic character oracid-base properties of the oxides

Prediction and Discovery of Germanium

1869-1871

Mendelejew Proposal of the “Periodic Table of Elements”, Prediction of properties of the undiscovered element 32 based on the periodicity concept “Table of Elements”

1886 Winkler Discovery of “Germanium” in a silver containing mineral

Properties

Proposed by Mendelejew

1871

Properties found by Winkler

1886

Present State

relative atom mass 72 72.32 72.61 colour dark grey grey white grey white density [g/cm³] 5.5 5.47 5.32 specific heat capacity [J/gK] 0.306 0.318 0.310 melting point [°C] high - 937.4 valency 4 4 4 oxide formula AO2 GeO2 GeO2 density [g/cm³] 4.7 4.703 4.228 acid/base

properties predominantly

acid acknowledged acknowledged

chloride formula ACl4 GeCl4 GeCl4 density [g/cm³] 1.9 1.887 1.8443 boiling point [°C] 60 - 100 86 83.1 ethyl compound formula A(C2H5)4 Ge(C2H5)4 Ge(C2H5)4 density [g/cm³] 0.96 0.99 0.991 boiling point [°C]

160 163 162.5

Historical Development of Understanding Chemical Bonds

1789 Lavoisier Theory of radicals

1807 1812

Davy Berzelius

Chemical bonds as electrochemical attraction, Discrimination between electropositve and electronegative elements

1852 Frankland Definition of “valence” as the ability of a given atom to form a compound with a defined number of other atoms (valency)

1857/58 Kekulé Kolbe Couper

Multiple carbon bonds in organic substances, first cyclic structure of benzene

1861 Butlerov Theory of chemical structures, determined by valence bonds

1874 van’t Hoff Le Bel

Stereochemistry

1910 Stark Falk Nelson

Coherence between valency and outer electrons (term “valence electrons”)

1916-19 Lewis Langmuir Kossel

Octet rule (noble gas shells), ionic and covalent bonds, covalent bonds as shared electron pairs

1927-29 Hund Mulliken Lennard-Jones

Quantum mechanical LCAO-MO-theory

1927-31 Heitler, London, Slater, Pauling

Quantum mechanical “valence bond theory”

1931 Pauling Hybridisation

Types of Chemical Bonds (1)

Octet rule:

The electron configuration of noble gases (s2, s2p6, s2p6d10, s2p6d10f14 – fully saturated shells) have the highest stability. Every atom tries to reach the electron configuration of the next neighboured noble gas by donating or accepting electrons. (8 valence electrons for elements of the 2nd and 3rd period) Please note: At the higher periods also other electron configurations, like (n-1)d10, (n-1)d10 (n)s2, ((n-1)d5(n)s2 can be preferred.

Covalent bonds

- sharing of electron pairs (electrons have different spins) between the bonded atoms

- If the partners are equal, the electron pair belongs to both partners in equal proportions, no dipole momentum can be observed.

- If the partners are different, the electron pair shifts to the atom with the stronger electron affinity (electron negativity). The bond will be polarised.

- dominates if difference of electron negativity is less than 1.7 - Valence Shell Electron Pair Repulsion Model (VSEPR):

Isolated electron pairs cause angled molecules (e.g. H2O).

CH4 NH3 H2O

109.5° 107° 104.9°


Ionic bonds

- Move of electrons from one partner to the another, ions electrically charged arise

- Bond is based on electric attraction of opposite ion charges. - dominates if difference of electron negativity is higher than 1.7

atomic bond polarised covalent bond ionic bond

polarisation of the covalent bond

polarity of the ions

There exists a continuum between covalent and ionic parts of bonds!

Molecule Ionic part of the bond

Molecule Ionic part of the bond

LiF 0.87 NO 0.015 LiCl 0.73 CO 0.01 LiBr 0.59 HCl 0.18 LiI 0.55 HBr 0.12

CsCl 0.75 HI 0.05 BaO 0.43 H2 0


Metallic bonds

- atom cores form a crystal lattice, valence electrons and orbitals are delocalised over the whole crystal (“electron gas”)

- exits only in solid or liquid metals - The energy difference between the “highest occupied molecule orbital”

(HOMO) and the “lowest un-occupied molecule orbital” (LUMO) is responsible for electrical conductivity: - low in case of metals (easy and fast electron transition), - moderate in case of semiconductor metals - high in case of isolators

Intermolecular interactions

- van der Waals attraction (weak interactions between the molecules, in general)

- Hydrogen bridging bonds § between acid H atoms and O, N or F atoms (2nd period) § intermolecular or intramolecular

Formic acid (intermolecular H bridging bounds)

Maleic acid (intramolecular

H bridging bounds)

Valency and Oxidation State Numbers

→ describe the number of electrons which one atom spends or attracts in a molecule → is the charge of an atom/ion, which would occur,

if the reaction considered is described as a heterolytic reaction forming ions

Oxidation states: - are 0 for the elements in general (also in molecules Ax, e.g. H2, O2, P4, S8) - are negative if a atom attracts electrons (corresponding to charge) e.g. O2-: -2, F-: -1 - are positive if a atom spends electrons e.g. Na+: +1, Fe3+ : +3 - within a molecule the sum of oxidation states must be 0 (condition of electroneutrality) - within an ion the sum of oxidation states must give the overall charge of the ion e.g. SO4

2-: S → +6, O → -2; 1 ⋅ (+6) + 4 ⋅ (-2) = -2 - within a chemical equation the sum of oxidation states must be equal on both sides e.g: 2 SO2 + O2 → 2 SO3 left side: 2 ⋅ (+4) + 4 ⋅ (-2) + 2 ⋅ (0)= 0 right side: 2 ⋅ (+6) + 6 ⋅ (-2) = 0 Mg + 2 H+ → Mg2+ + H2 left side: 1 ⋅ (0) + 2 ⋅ (+1) = 2 right side: 1 ⋅ (+2) + 2 ⋅ (0) = 0 Valency state numbers:

- are the absolute (positive values) of oxidation state numbers e.g. Na+: I, O2-: II

- are written in Roman numerals

Quantum Mechanical Concepts of Molecular Bonds 1. Theory of Molecular Orbitals (MO Theory)

Forming a molecule the atoms have to overlap their atom orbitals. → “Linear combination of atom orbitals to molecular orbitals”

(LCAO-MO theory) by Hund, Mulliken, Lennard-Jones (1927-1929)

positive interference

negative interference

no interference

Algebraic signs are related to the angular part of the wave function, not to a charge!

- interference can occur, if the atom orbitals have the same symmetric properties with respect to the bond axis

- number of MO is equal to the number of interacting atom orbitals - positive interference: bonded MO, decrease of energy - negative interference: anti-bonded MO, increase of orbital energy

σ orbital

π orbital

δ orbital

- number of bonds = number of bonded MO - number of anti-bonded MO

Quantum Mechanical Concepts of Molecular Bonds 2. Theory of Valence Bonds (VB Theory)

Heitler, London, Slater, Pauling (1927-1931) Coupling of unpaired electrons to bonds gives molecular valence structures.

H• + H• → H-H H? H? H? H?

The coupled electron pair can belong to - both partners: covalent electron pair - one atom: ionic electron pair

Favoured valence structure: - maximized number of covalent bonds - structures with short bond lengths - ionic structures, where the electron pair is situated at the atom with

highest electron affinity - ionic structures, where opposite charges are situated in the next

neighbourhood

The overall wave function is represented by the linear combination of all possible valence structures. Each valence structure can be transformed easily to another valence structure (resonance). Valence structures are mesomeric borderline cases of the reality.

Quantum Mechanical Concepts of Molecular Bonds 3. Hybridisation

Linus Pauling (1931)

→ Linear combination of s, p (and d) orbitals forms new hybrid orbitals. → Combination of LCAO-MO method and VB theory

Overlaying the atom orbitals

Resulting hybrid orbitals

Type of hybrid orbital

Involved atom orbitals

Geometric form

Type of the molecule

Geometry of the molecule

Example

sp s, px linear AB2 linear BeCl2 sp² s, px, py triangle AB3

AB2 triangle V form

BF3 SO2

sp³ s, px, py, pz tetrahedral AB4 AB3

AB2

tetrahedral triangle pyramid V form

CH4 NH3

H2O sp²d s, px, py, dxy quadratic AB4 quadratic XeF4 sp³d s, px, py, pz,

2zd

triangle bipyramidal

(2 tetra-hydrons)

AB5

AB3

AB2

triangle bipyramidal T form linear

PF5

ClF3

XeF2 sp³d s, px, py, pz,

22 yxd

− quadratic pyramid

AB5 quadratic pyramid

BrF5

sp³d² s, px, py, pz,

2zd , 22 yx

d−

octahedron AB6 AB5

AB4

octahedron quadratic pyramid quadratic

SF6 BrF5

XeF4

Special Cases of Hybrid Orbitals

Ethylene

(C-C double bonds)

= σ bond + π bond (sp2 hybrid orbitals)

Acetylene

(C-C triple bonds)

= σ bond + 2 π bonds (sp hybrid orbitals)

Diborane B2H6

(“electron shortage compounds”)

2 electrons triple center bond

Benzene

(“aromatic systems”)

delocalised conjugated π system

Multiple bonds occur only with elements of the 2nd period. At higher periods they will be “prevented” by polymerisation (e.g. CO2 vs. SiO2).

Literature/References for Figures

(1) Arnold Frederik Holleman, Egon Wiberg, Lehrbuch der anorganischen Chemie

101st edition, Berlin [u.a.] : de Gruyter, 1995 A lot of pages (2033), and a lot of detailed information, the standard book for inorganic chemistry in Germany

(2) Gisbert Großmann, Jürgen Fabian, Lehrwerk Chemie, Lehrbuch 1 „Struktur und Bindung – Atome und Moleküle“, 6th edition, VEB Deutscher Verlag für Grundstoffindustrie, 1989 The first book from a series, related to all topics of chemistry studies. Small and compact (252 pages). It was used as the standard book in the former GDR.

(3) P.W. Atkins, Physical Chemistry, 6th edition, Oxford University Press, 1998 A well readable book on basic level about all topics on physical chemistry.

(4) Richard Stephen Berry, Stuart A. Rice, John Ross, Physical chemistry, 2nd edition, Oxford Univ. Press, 2000

Trends in the Periodic Table of Elements



Lanthanides

Actinides

acid strenght of oxides

basic strenght of oxides

non-metallic character

metallic character

electron affinity/negativity, ionisation energy

ability to be oxidised

electron affinity, ionisation energy,non-metallic character,acid strenght of oxides,oxidation state (valency)



Lanthanides

Actinides



Lanthanides

Actinides

acid strenght of oxides

basic strenght of oxides

non-metallic character

metallic character

electron affinity/negativity, ionisation energy

ability to be oxidised

electron affinity, ionisation energy,non-metallic character,acid strenght of oxides,oxidation state (valency)

Course on Inorganic Chemistry

Chapter 2

Chemical Reactions

The Chemical Equilibrium

Consider the reaction α A + β B → 1k γ C + δ D with r1 = k1 * [A]α * [B]β

If back reaction

γ C + δ D → −1k α A + β B with r-1 = k-1 * [C]γ * [D]δ also occurs, we have a chemical equilibrium described by

βα

δγ

− ⋅⋅

==]B[]A[]D[]C[

kkK

1

1

“Mass Action Law” (K – equilibrium constant, k1 and k-1 – rate constants for the reactions,

[A], [B], [C], [D] – concentrations or partial pressures, α, β, γ, δ – reaction orders ) Transition state theory (Eyring)

reaction coordinate

catalystener

gy

A + B

C + D

activated complex

reaction coordinate

catalystener

gy

A + B

C + D

activated complex

k1 = k0, 1 * exp (-EA, 1/RT)

k-1 = k0, -1 * exp (-EA, -1/RT)

EA, 1 ≠ EA, -1

In case of equilibrium r1 = r-1 ≠ 0

→ dynamic equilibrium

Special cases:

(1) Nernst’s distribution law 2phase,A

1phase,A

cc

K =

(2) Henry Dalton’s law phasegas,A

solutionliqiuid,A

pc

RTK

K ==′

(3) electrolytic dissociation ]AB[

]A[]B[Kc

−+ ⋅=

(Kc < 10-4 - weak electrolytes, Kc > 10-4 - intermediate electrolytes, Kc → 8 - strong electrolytes (full dissociation))

Le Chatelier’s Principle (1888)

A system in equilibrium, when subjected to a disturbance, responds in a way that trends to minimise the effect of disturbance. (1) Increase of temperature

→ favours the endothermic reaction Decrease of temperature → favours the exothermic reaction

(2) Increase of pressure → favours the reaction with ∆rV < 0 Decrease of pressure → favours the reaction with ∆rV > 0

(3) Increase of the concentration of one reactant → favours the reaction consuming this reactant Removal of one reactant → favours the reaction of its re-formation

Note: Catalysts increase both reaction rates r1 and r-1, so that the equilibrium is reached faster, but under identical reaction conditions the distribution between the reactants doesn’t change.

Reduction and Oxidation

Reducing agent

Oxidising agent + electrons

Oxidation

ReductionReducing

agentOxidising

agent + electronsOxidation

Reduction

Oxidation number/oxidation degree:

charge of an atom, which would occur, if the reaction considered is described as a heterolytic reaction forming ions Examples: elements ±0 HCl Oxidation number of hydrogen +1 Oxidation number of chlorine – 1 H2O Oxidation number of hydrogen +1 Oxidation number of oxygen – 2 The negative charge must attributed to the partner with the highest electron negativity (see Periodic Table of Elements!!).

Electrochemical Potentials

electrical connection

membrane

ZnSO4 solution CuSO4 solution

Zn pole Cu pole

electrical connection

membrane

ZnSO4 solution CuSO4 solution

Zn pole Cu pole

Galvanic cell (voluntary) Anode (negative pole - oxidation):

Zn → Zn2+ + 2 e-

Cathode (positive pole - reduction): Cu2+ + 2 e- → Cu

The back reaction is “electrolysis” forced by applying the opposite voltage.

Electrochemical Potential Series

- Potentials are relative values.

→ Normalisation on H2/2 H+ standard electrode (= ± 0.000 V) - Nomenclature: reduced/oxidised species (Na/Na+, 2 Cl-/Cl2) - low potential (negative – e.g. alkali metals)

= high reduction power = easy to be oxidised - high potential (positive – e.g. noble metal cations)

= high oxidation power = easy to be reduced

→ allow to predict reactions (∆G = Z*F*ε) → applied in practice in electrochemical processes (e.g. galvanisation), in batteries

and fuel cells

Concentration dependency of potentials

Nernst Equation: .dRe

.Ox0 c

clgFZTR ⋅

⋅⋅+ε=ε

ε – potential ε0 – standard potential (see tables)

R – gas constant T – temperature

Z – number of electrons, which should be donated or accepted F – Faraday constant

cOx./cRed. – concentration of oxidised/reduced reactants (like in the mass action law)

Setting ε to 0, it is possible to get the equilibrium constant K.

Normalised potentials for acid (pH = 0) and basic (pH = 14) solutions (at 25 °C)

a) metals acid solution basic solutionacid solution basic solution

b) non-metallic elements and compounds acid solution basic solutionacid solution basic solution

→ Power of oxidising agents which are reduced increases in acid solutions. Power of reducing agents which are oxidised increases in basic solutions.

The Acid-Base Concept Proposed by Brönstedt and Lowry

- acids = proton donators, bases = proton acceptors - valid for water and other protical solvents (e.g. liquid NH3) - acid reaction: HX + H2O X- + H3O

+ (H3O

+ - oxonium ion, which will be solvatisated, hydronium ion = [H3O ⋅ 3 H2O]+)

- base reaction: M-OH M+ OH-

- autoprotolysis reaction of water: 2 H2O OH- + H3O+

K = 10-14, pH = -log [H3O+]

Acid anhydrides = compounds (oxides or metal cations) forming Brönstedt acids first by the reaction with water e.g. SO3 + 2 H2O H2SO4 + H2O HSO4

- + H3O+

Al3+ + 7 H2O Al(OH2)63++ H2O [Al(OH2)5(OH)]2++ H3O

+

Amphoteric compounds (ampholytes) Compounds (mostly oxides), which can form acid and base ions:

Al(OH)3 + 3 H3O

+ [Al(OH2)6]3+ (pH < 5) Al(OH)3 + 3 OH- [Al(OH)6]3- (pH > 9)

Between pH 5 and 9 solid Al(OH)3 falls out.

The Acid-Base Concept Proposed by Lewis (1923)

- acids = electron pair acceptors, bases = electron pair donators

Lewis acid + Lewis base Lewis acid-base complex Lewis acids: cations or electron shortage compounds,

which can attract electron pairs BF3, AlH3, SO3, H

+, Fe2+ Lewis bases: anions or compounds with unbounded electron pairs

F-, H2O, OH-, NH3, CN- → Lewis acid-base concept includes partially redox reactions.

Principle of hard and soft acids and bases (HSAB principle - by Pearson 1963) Stability of the acid-base complex is high if there react hard acids with hard bases or weak acids with weak bases. hard acids: cations with small diameters, high positive charge and no non-bonded

electrons, → H+, cations from s1, s2, s2p1 and d10s2p2 elements

→ forming mainly ionic bonds soft acids: cations with large diameters, low positive charge and non-bonded

electrons, → cations from transition metals with d10s2 configuration (type B cations)

→ forming mainly covalent bonds

hard bases: anions with a central atom highly charged and possessing a high electronegativity

soft bases: anions with a central atom low charged and possessing a low electronegativity

(hard) Anions of F > O >> N, Cl >Br, H >S, C > I > P (weak) Hard or soft properties of Lewis acids and bases can be found only experimentally. Additionally strength of Lewis acids and bases must be considered! Strong acids + strong bases give stable complexes every time (H+ + H- → H2), but selectivity is influenced by hard or soft character (Al 2S3 + HgO → Al2O3 + HgS).



101st edition, Berlin [u.a.] : de Gruyter, 1995

(3) Gisbert Großmann, Jürgen Fabian, Lehrwerk Chemie, Lehrbuch 2 „Struktur und Bindung – Aggregierte Systeme und Stoffsystematik“, 5th edition, VEB Deutscher Verlag für Grundstoffindustrie, 1987

(3) P.W. Atkins, Physikalische Chemie, 2nd reprint of 1st edition, VCH Verlagsgesellschaft Weinheim, 1990


Chapter 3

Noble Gases

Overview About the Group

Group Members

Helium (He)

Neon (Ne)

Argon (Ar)

Krypton (Kr)

Xenon (Xe)

Radon (Rn)

Atom Number

2 10 18 36 54 86

Rel. Atomic Mass

4.00 20.18 39.95 83.80 131.29 [222] (radio-active)

Discovery 1895 Ramsey

1898 Ramsey

1894 Ramsey, Rayleigh

1898 Ramsey

1898 Ramsey

1900 Dorn,

Rutherford, Soddy

Percentage in air [Vol.-%]

0.000524 0.00182 0.9340 0.000114 0.000087 6 * 10-18

Electron configuration: s²p6(d10) – fully saturated electron shells → very poor or no reactivity Industrial manufacturing: Rectification of air (Linde process) Use: Inert gas in lamps and in high temperature processes

(Ne, Ar, Xe) Balloon gas (He) Medicine (Rn as source of α42 species)

Physical and Chemical Properties

Group Members

Helium (He)

Neon (Ne)

Argon (Ar)

Krypton (Kr)

Xenon (Xe)

Radon (Rn)

melting point [°C]

-272.1 (2.5 MPa)

-248.6 -189.4 -156.6 -111.5 -71

boiling point [°C]

-268.9 (4He)

-246.0 -185.9 -152.9 -107.1 -61.8

vaporization enthalpy [kJ/mol]

0.092 1.86 6.28 9.68 13.70 18.02

1st ionisation energy [eV]

24.58 21.56 15.76 14.00 12.13 10.7

Low Temperature Properties of Helium

- lowest boiling and melting temperature of all substances - cannot be frozen under atmospheric pressure (this needs 25.5 bars) - Helium I (normal fluid) and Helium II (super fluid)

He(I) → He(II) at -270.97 °C (2.18 K)/1 bar for 4He first at extreme low temperatures for 3He

- different physical properties of 3He and 4He boiling points: 3.20/4.21 K density: 0.08/0,14 g/cm³

→ easy separation of isotopes possible Ionisation potential of highest reactive elements:

O2 - 12.75 eV, similar to Xe F2 - 17.4 eV, higher than Kr and Xe Cl2: - 12.9 eV, similar to Xe Br2: - 11.76 eV

First noble gas compound: - “clathrates” (“enclosed compounds”, “cage compounds”) - XePtF6 by Barlett (1962, theoretically predicted by Pauling 1933) Known noble gas compounds - RnF2, fluorides, oxides and oxifluorides of Xe, chlorides of Xe, KrF2 - no compounds of He, Ne , Ar

The Air Rectification Process by Linde

Joule Thomson effect: Gases can be cooled by adiabatic expansion, if

temperature δa is less than inversion temperature and µJT (Joule Thomson coefficient) is positive.

The Linde process

Joule Thomson parameter and inversion temperature for

different gases

(heating)

(cooling)

(as the „ideal gas“)

(heating)

(cooling)

(as the „ideal gas“)

cooler

compressor

air inlet

liqiud air

compressed air

expanded air

throttle valve

cross flowheat exchanger

δA, pA

δE, pE

cooler

compressor

air inlet

liqiud air

compressed air

expanded air

throttle valve

cross flowheat exchanger

δA, pA

δE, pE

Process scheme: 1. Air in compressed to 200 bar (pA) 2. Compressed air is cooled to remove compression heat 3. Expanding of cooled compressed air followed 4. Expanded air cools compressed air 5. Air is compressed again (like 1.)

Cooling effect of Linde process:

δA - δE = µJT * (pA – pE) *(273.15/(273.15 + δA))² Fractions of the technical rectification

→ further purification in additional rectification steps

Halogen Compounds of Noble Gases

Xe + F2 → ° tubeNiC ,400 XeF2

(colourless solid)

Xe + 2 F2 → =° 5:1/,6.0,400 2FXeMPaC XeF4 (colourless solid)

Xe + 3F2 → =°− 20:1/,5,250200 2FXeMPaC XeF6

(colourless solid)

Kr + F2 → °− mbarC 20,183 KrF2 (colourless solid)

Molecular structures of XeF2, XeF4 and XeF6

- reaction is possible after activation of fluorine (F2 → 2 F) by heat, UV radiation, microwaves, electrical discharges or radiation

- stability: - increases with increasing atomic number of noble gas

atom (RnF2/XeF2 (stable) >> KrF2 (stable until – 70 °C) > ArF2 (not reported)) - decreases with increasing atomic number of halogen atom (XeF2 (stable) >> XeCl2 (unstable)> XeBr2 (unstable)) - decreases with increasing oxidation state of noble gas atom (XeF2 > XeF4 > XeF6 (all stable, but increasing formation enthalpy +164/+278/+361 kJ/mol), XeF8 (not reported))

- all noble gas halogen compounds have strong oxidation power § XeF2: Cl- → Cl2, IO3

- → IO4, BrO3- → BrO4

- (all in aqueous solutions), Fluorination of NO2 to FNO2, reaction with F2 to XeF4 and XeF6

§ XeF4: Pt → PtF4, Hg → Hg2F2 § XeF6: Hg → HgF2, AuF3 → Au(V) § KrF2: ClF3 → ClF5, Ag → AgF2, Hg → HgF2,

[KrF]+ strongest known oxidation agent 7 KrF2 + 2 Au → 2 [KrF][AuF6] → AuF5 + Kr + F2

Oxygen Containing Compounds of Noble Gases

→ only known compounds: XeO3, XeO4, H4XeO6, XeOF2, XeO2F2, XeOF4, XeO3F2, XeO2F4

Molecular structures of XeO3 and XeO4

Xenon(VI)-oxide (XeO3)

- preparation: XeF6 + 3 H2O → XeO3 + 6 HF 3 XeF4 + 6 H2O → Xe + XeO3 + 12 HF - properties: colourless crystals,

soluble in water (> 1 mol/l), weak acid (pKs = 10.5) high oxidation power (Cl- → Cl2, Mn (II) → Mn (IV)) explosive

Xenon(VIII)-oxide (XeO4)

- preparation: basic hydrolysis of XeO3

XeO3 + OH- → HXeO4-

2 HXeO4- + 2 OH- → XeO4

6-+ Xe + O2+2 H2O - properties: XeO4 – yellow liquid (< - 40 °C)/colourless gas,

XeO46- yellow solutions

XeO4 – explosive above – 40 °C strong oxidation power (ClO 3

- → ClO4-, Cr3+ → Cr2O7

-, Mn2+ → MnO4

-,(IO3- → IO4

-)

Oxiflouride Compounds - preparation: deep temperature hydrolysis of XeF4, reaction of xenon fluorides with xenon oxides - properties: colourless crystals, which can be hydrolysed, poor stability





(3) P.W. Atkins,

Physikalische Chemie, 2nd reprint of 1st edition, VCH Verlagsgesellschaft Weinheim, 1990


Chapter 4

Hydrogen

Overview

- discovered 1766 by Cavendish - lightest element - third most common element by atom percentage,

ninth most common element by mass percentage - occurs in nature mostly as oxide (water H2O)

Hydrogen isotopes Name protium H deuterium D tritium T atom core composition

1 proton 1 proton + 1 neutron

1 proton + 2 neutrons

rel. atomic mass

1.0078 2.0141 3.0160

natural percentage

99.9855 % 0.0145 % 10-15 %

Electron configuration: s1 → needs to spent or to accept one electron → occurs in elementary form as diatomic H2 ortho and para hydrogen

spins of protons

ortho-hydrogen para-hydrogenatom cores

electron shell

ortho-hydrogen para-hydrogenatom cores

electron shell

ortho-hydrogen

para-hydrogen

perc

enta

ge o

fpar

a-hy

drog

en [%

]

perc

enta

ge o

forth

o-hy

drog

en [%

]

absolute temperature

ortho-hydrogen

para-hydrogen

perc

enta

ge o

fpar

a-hy

drog

en [%

]

perc

enta

ge o

forth

o-hy

drog

en [%

]

absolute temperature

- o-H2 p-H2 + 0.08 kJ/mol - ratio at 25 °C: 75/25 - separation by adsorption on

alumina at 20.4 K and 50 mbar - differences in physical properties

(melting and boiling points, cp, vapour pressures)

Chemical properties Homolytic dissociation energy (H2 → 2 H): 436.2 kJ/mol

→ catalytic activation by high dispersed transition metals (e.g. Pt, Pd)

Heterolytic dissociation energy (H2 → H+ + H-): 1675 kJ/mol Oxidation enthalpy (2 H2 + O2 → 2 H2O): -572.04 kJ/mol Reduction enthalpy (Ca + H2 → CaH2): -184 kJ/mol

Manufacturing and Use of Elementary Hydrogen

Industrial manufacturing: world production 35 mill. tons/year (1990)

Steam cracking/Steam reforming of oil and natural gas (>90 %) CH4 + H2O CO + 3 H2 (700-830 °C, 40 bar, Ni catalyst) Coal gasification C + H2O CO2 + H2

Water shift reaction CO + H2O CO2 + H2

Synthesis gas is a mixture of CO and H2 (traces of CO2, and H2O)

Chlorine alkali electrolysis NaCl + H2O → NaOH + 1/2 Cl2 + 1/2 H2

Laboratory manufacturing: Reaction of non-noble metals (Zn, Ca, Mg)

with diluted acids (HCl, H2SO4, HNO3) M + 2 H+ → M2+ + 2 H → M2+ +H2

(2 H = “status nascendi” = high reactive atomic hydrogen) Reaction of metallic Al or Si with hot NaOH giving aluminates and silicates and H2

Use: Ammonia production

(Haber Bosch process, N2 + 3 H2 2 NH3) Organic chemistry (Hydrogenation, reducing agent)

Inorganic chemistry (HCl synthesis, reducing agent for metal manufacturing) Food industry (fatty acid hydrogenation) Fuel Cell Technology

Binary Hydrogen Compounds – Ionic Compounds General: - formed with elements of the 1st and 2nd main group by direct synthesis from the

elements

- nomenclature: [Name of the metal] – ([number of H atoms]) – hydride e.g. Magnesium(di)hydride – MgH2

Structure and Properties: - bond between the metal and the hydrogen is primarily ionic

- metal ions possess positive charge, hydrogen is charged negatively - exothermic compounds with salt crystal structures, high decomposition

temperatures (300-1000 °C), electrical conductivity in molten state

- solution in water under decomposition to hydroxides and hydrogen - strong reducing agents, industrial use for manufacturing pure elements (e.g. LiH,

NaH, CaH2)

- exception: BeH2 is a typical covalent compound

Reactions: with halogens to metal halogenides + hydrogen e.g. CaH2 + X2 → MeX2 + H2

with oxygen to oxides and water e.g. CaH2 + O2 → CaO + H2O (500 °C) with nitrogen to nitrides and hydrogen e.g. 3 CaH2 + N2 → Ca3N2 + 3 H2 (500 °C) with carbon to carbides and hydrogen e.g. CaH2 + 2 C → CaC2 + H2 (>700 °C)

Binary Hydrogen Compounds – Metallic Compounds (1) General: - formed with transition metals and the metals of the III.-VI. main group Preparation:

- by direct synthesis from the elements giving non-stochiometric compounds

hydrogen uptake (mol H/mol metal)

hydride phase

mixed phase(solution + hydride)

plateau regionhy

drog

en p

ress

ure

solutionphase

hydrogen uptake (mol H/mol metal)

hydride phase

mixed phase(solution + hydride)

plateau regionhy

drog

en p

ress

ure

solutionphase

- by reaction of halogenides with LiH, NaBH2 or LiAlH4

EHaln + n H- → EHn + n Hal-

Stability:

- IIIb and IVb groups: exothermic compounds, stable at room temperature - Vb group and CrH: endothermic compounds, meta-stable - VIb-VIIIb groups: very unstable or not discovered - Ib, IIb and IIIa-VIa groups: endothermic compounds, stable only at low

temperatures - stability decreases with increasing hydrogen content

(VH stable at room temperature, VH2 decomposes)

Group IIIb IVb Vb VIb VIIb VIIIb Ib IIb Hydrogen metal ratio x of the compound EHx

=31 =2 =2 =22 n.d. =22 1 2

Stability increase Hydrogenation catalysts

1 – including lanthanoides and actinoides 2 – only known from Cr, Ni (at high pressures) and Pd n.d. – not discovered

Binary Hydrogen Compounds – Metallic Compounds (2) Structure and Physical Properties:

- “inlay compounds” – no changes in the metal lattice structure

- hydrogen atoms occupy lattice gaps, they can move inside the gap

- presence of cationic and anionic hydrogen

- conductors using free electrons

- in gas phase linear molecules H—M—H

Reactions and Use:

- reaction with water under decomposition to hydroxides and hydrogen - manufacturing of high purity metals - hydrogen storage

Binary Hydrogen Compounds – Covalent Compounds (1) General: - formed with non-metallic elements of the III.-VII. main group - high industrial importance Preparation: - by direct synthesis from the elements (e.g. Haber Bosch process for ammonia) - reaction of metal compounds of elements of the III.-VII. main group with acids (e.g.

CaF2 + H2SO4 → CaSO4 + 2 HF – industrial process) Structure:

planar triangle(IIIa group elements)

tetrahedron(IVa group elements)

pyramidal(Va group elements)

planar angled(VIa group elements)

planar triangle(IIIa group elements)

tetrahedron(IVa group elements)

pyramidal(Va group elements)

planar angled(VIa group elements)

- hydrogen possesses positive charge for H-Hal, H2O-H2Se, NH3, CH4 - multiple centred bonds (coordination number of H = 2, equal bond length) in Be-H

and B-H compounds via anionic hydrogen bridging bonds (polymerisation)

- association of hydrides from elements of the 2nd period

via cationic hydrogen bridging bonds (longer than covalent bonds)

(HF)x (solid)

(HF)6 (gaseous)

Binary Hydrogen Compounds – Covalent Compounds (2) Physical and Chemical Properties:

- non-conductors - high volatility except hydrocarbon compounds from N, O and F (much higher

melting and boiling points because of cationic hydrogen bridging bounds) and from B (dimerisation)

melting points boiling points

periodperiod

melting points boiling points

periodperiod

- exothermic and stable compounds (without higher periods) - solubility in water

o H-Hal: high solubility with strong acid reaction o H2X (VIa group): high solubility with weak acid reaction o NxHy: high solubility with strong basic reaction o H3X (Va group since P): low solubility with low basic reaction o H4X (IVa group): no solubility o (H3B)x: no solubility → dissociation 2 EHn EHn+1

+ + EHn-1-

- well soluble in ethers - use as polar solvents: H2O, NH3 (liquid), HF (liquid)

use as non-polar solvents: higher hydrocarbons (C = 6…12)

- reducing power (EHn + (n+p)/2 X2 → n HX2 + EXp) F2 > O2 > Cl2 >Br2 …, correlates to electronegativity and to normalised electrochemical potentials

Binary Hydrogen Compounds – Covalent Compounds (3) Higher Hydrogen Compounds

- molecules with more than one single or multiple bonded element atoms (especially with elements of 2nd period)

Reactions

- protonation/deprotonation (Va-VIIIa group elements) H+ + H2O H3O

+ NH3 + H3O

+ NH4+ + H2O

NH3 NH2- + H+

- accepting hydride ions (IIIa group elements):

BH3 + H- BH4-

AlH3 + H- AlH4-

Heavy and Super-heavy Water

H2O

“light water” D2O

“heavy water” T2O

“super-heavy water”

rel. molecular mass 18.02 20.03 22.03 density (25 °C) [g/cm³] 0.997 1.104 1.214 maximum density [g/cm³] / Temperature of density maximum [°C]

1.000/3.98 1.106/11.23 1.215/13.4

melting point [°C] 0.000 3.81 4.48 boiling point [°C] 100.00 101.42 101.51 dissociation constant pKW (25 °C)

14.000 14.869 15.215

Toxicity low (salt-free) high radioactive Industrial manufacturing : Electrolysis of used technical electrolyte

solutions → enrichment of D2 during the end of the process because of lower reaction rate

Use: Nuclear industry, Studies on reaction mechanisms

(H-D exchange)





(3) P.W. Atkins,

Physikalische Chemie, 2nd reprint of 1st edition, VCH Verlagsgesellschaft Weinheim, 1990


Chapter 5

Halogens


Group Members

Fluorine (F)

Chlorine (Cl)

Bromine (Br)

Iodine (I)

Astatine (At)

Atom Number 9 17 35 53 85 Rel. Atomic Mass

19.00 35.45 79.9 126.90 209.99

Discovery 1886 Moissan

1774 Scheele

1826 Balard

1811 Courtois

1940 Corson,

McKenzie, Segré

Percentage on earth [Mass-%]

0.06 0.11 6 * 10-4 5 * 10-5 3 * 10-24 (radio-active)

melting point [°C]

-219.62 -101.00 -7.25

113.60 300

boiling point [°C]

-188.13 -34.06 58.78 185.24 335

state at room temperature (25 °C) and 1 bar

colourless - weak yellow gas

yellow-green gas

dark brown liquid

violet crystals

with metallic brilliance

solid

Electron negativity

4.0 3.0 2.8 2.5 2.2

valence numbers in compounds

-1 -1…+7 -1…+7 -1…+7 -1…+7

Reducing Power Oxidation power

Electron configuration: s²p5(d10) – need of accepting one electron

or loosing 7 electrons for full saturation of electron shells

→ very high up to extreme reactivity → occurs in nature only in compounds → in gas phase diatomic molecules X2

Manufacturing and Use of Elementary Fluorine

→ Fluorine is the elements with the highest electronegativity (4.0). It is the element with the highest reactivity. Elementary fluorine cannot be formed by any chemical reaction.

Natural Sources: fluorspar – CaF2 (main source, 5 * 106 t/year) fluorapatite – 3 Ca3(PO4)2 · CaF2 (with 2…4 mass-% F) cryolite – Na3AlF6

Manufacturing in industry: 1. Conversion of fluorspar to hydrofluoric acid

CaF2 + H2SO4 → CaSO4 + 2 HF 2. Electrolysis of hydrofluoric acid to fluorine in water-free molten KF · 2 HF

(melting temperature: 72 °C) 2 HF → H2 + F2

Process data

voltage: current: current density: temperature: yield

8-12 V 4-15 kA

0.5-0.15 A/cm² 70-130 °C

90-95 %

(relative to the current consumed)

3. Purification of F2 by freezing out un-reacted HF at -100 °C

Properties and Application:

- high toxic in elementary form (essentially as ionic fluoride) - one of the strongest oxidation agents (H2O + F2 → 2 HF + 0.5 O2) - heavy reaction with most of the other elements

even at room temperature (except He, Ne, Ar) - used for industrial synthesis of UF6, SF6, CF4 and fluorographite (electrodes in

batteries) - surface fluoridation (Teflon)

Manufacturing and Use of Elementary Chlorine

Natural Sources: sodium chloride – NaCl

(main source, from mining or from seawater, 170 * 106 t/year, purification and enrichment up to 99 %)

potassium chloride – KCl (mostly used as fertilizer) other natural salts: KMgCl3 · 6 H2O,

MgCl2 · 6 H2O, KMgCl(SO4) · 3 H2O

Manufacturing in industry:

1. Electrolysis of NaCl brines (2 H2O + 2 NaCl → H2 + 2 NaOH + Cl2) Restrictions to the process:

- prevention of formation of hypochlorite in the solution (2 OH- + Cl2 → OCl- + Cl-) - suppressing of contact between H2 and Cl2 (→ 2 HCl – danger of explosions)

- mercury process (40 % of world production) - diaphragm process (40 % of world production) - membrane process(20 % of world production)

2. Electrolysis of concentrated hydrochloric acid 2 H+ + 2 Cl- → H2 + Cl2 3. Thermal oxidation of hydrogen chloride (Deacon process) 4 HCl + O2 2 H2O + 2 HCl + 2 Cl2

catalyst: CuCl2 (Deacon process – 350 °C) or MnO2 (Weldon process)

Properties:

- yellow green, suffocative smelling gas - soluble in water (0.0921 mol /l = 6.6 g/l) - high toxic in elementary form (essentially as ionic chloride) - high reactivity, especially with non-noble metals and hydrogen

(but less than fluorine) - reactivity is increased by adding small amounts of water

(forming of traces of ClO- initiators) - high oxidation power (less than fluorine)

Application:

- synthesis of organic chemicals (mainly vinyl chloride) - leaching agent in paper and pulp industry - inorganic chemicals, water treatment, cleaning and sanitation

The Mercury Process for Manufacturing Chlorine

General: - separation of chloride oxidation and hydrogen reduction - Step 1: Electrolysis of NaCl gives sodium solved in mercury (amalgam) and

gaseous chlorine. Anode: Cl- → 0.5 Cl2 + e- (ε = 1.24 V) Cathode: x Hg + Na+ + e- → NaHgx (ε = -1.66 V)

- Step 2: Decomposition of amalgam to Hg (recycling), NaOH and H2 NaHgx +H2O → 0.5 H2 + NaOH + x Hg

Process parameters

cell voltage:

current density: temperature:

NaCl concentration start: end:

electrochemical yield:

4.2 V 8-15 kA/m²

80 °C

310 g/l 260-280 g/l

94-97 %

Advantages: - pure 50 % sodium hydroxide solution without evaporation - high purity chlorine gas

Disadvantages: - need of higher voltage and energy compared to the diaphragm process - stronger brine purification requirements - care on preventing emissions of mercury

Halogen Oxygen Compounds – The Complete Reaction Network

Example: Chlorine

oxidation stateoxidation state

The Diaphragm Process for Manufacturing Chlorine

General: - separation of chloride oxidation and hydrogen reduction

by an asbestos membrane - Reactions:

Anode: Cl- → 0.5 Cl2 + e- (ε0= 1.36 V) Na+ + OH- NaOH Cathode: H2O + e- → H2 + OH- (ε0 = 0 (pH = 0)/-0.828 V (pH = 14)) NaCl Na+ + Cl-

Process parameters

cell voltage:

current density: final NaCl concentration: final NaOH concentration:

3.0-4.15 V 2.2-2.7 kA/m²

170 g/l 12-16 %

Diaphragm functionalities: - hindering of gas transport between the chambers - suppressing of contact between

H2 and Cl2 (but permeability for dissolved Cl2) - hindering of OH- transport to the anode

Advantages: - less requirements to NaCl purity - lower voltage and energy consumption

Disadvantages: - need of additional separation steps for NaOH and NaCl,

and of an evaporation step to enrich NaOH - oxygen content in the chlorine - care on preventing emissions of asbestos

The Membrane Process for Manufacturing Chlorine

General: - separation of chloride oxidation and hydrogen reduction by a Nafion membrane - Reactions:

Anode: Cl- → 0.5 Cl2 + e- (ε0= 1.36 V) Na+ + OH- NaOH Cathode: H2O + e- → H2 + OH- (ε0 = 0 (pH = 0)/-0.828 V (pH = 14)) NaCl Na+ + Cl-

Process parameters

cell voltage: current density:

final NaOH concentration: current yield:

3.15 V

2-3 kA/m² 35 % 95 %

with respect to NaOH

membrane materialsmembrane materials

Properties of the membrane: - thickness: 0.2 mm - ion-conductible, but non-permeable for the brine

Advantages: - pure NaOH without NaCl impurities - lower voltage and energy consumption than the mercury process - no use of mercury or asbestos, ecological most favoured process

Disadvantages: - high purity requirements to NaCl - low final concentration NaOH

and of an evaporation step to enrich NaOH - oxygen content in the chlorine - high costs and short lifetime of membranes

Manufacturing and Use of Elementary Bromine

Natural Sources: seawater (main source) residual solutions from potash (K2CO3) industry


Chlorine extraction of bromide ion containing brines (500000 t/a) (2 Br- + 2 Cl2 → Br2 + 2 Cl-)

- “Cold debromination” 1. acidification of seawater to pH = 3.5

with sulfuric acid (H2SO4) 2. extraction of formed bromine by “blowing out” with air 3. purification by adsorption with soda solution (a) and desorption with

H2SO4 and steam (b) 3 Br2 + 6 OH- → 5 Br- + BrO3

- + 3 H2O (a) 5 Br- + BrO3

- + 6 H+ → 3 Br2 + 6 H2O (b) - “Hot debromination” (major process)

1. Counter-current extraction of brines with a mixture of steam and Cl2 at 80 °C

2. Condensation of the steam containing Br2, Cl2 and H2O 3. Purification by distillation

Properties:

- brown high volatile liquid (melting point: -7.25 °C, boiling point 58.78 °C)

- soluble in water (0.2141 mol /l = 34.2 g/l) - high toxic in elementary form - quite high reactivity, less than fluorine and chlorine - reactivity is increased by adding small amounts of water

(forming of traces of BrO- initiators) - high oxidation power (less than fluorine and chlorine)

Application:

- synthesis of organic chemicals (mainly for medicine) - manufacturing of flame retardants (in decrease) - inorganic chemicals

Manufacturing and Use of Elementary Iodine

Natural Sources: - occurs in nature only in small concentrations

as iodide (I-) or iodate (IO3-)

- industrial sources: residual solutions from Chilean niter (NaNO3) industry (main source, containing mainly IO3

-), brines from crude oil and natural gas production


1. from residual solutions of Chilean niter production (50 %) - acidification of brines with H2SO3

(treatment with gaseous SO2) – reduction of IO3- to I-

(HIO3 + 3 H2SO3 → HI + 3 H2SO4) - comproportionation of iodine hydrogen with further

iodine acid (5 HI + HIO3 → 3 I2 + 3 H2O)

- purification by sublimation of the crude iodine 2. from brines from crude oil and natural gas production (50 %)

- similar process like for bromination extraction with Cl2/H2SO4 → “blow out” of iodine with air

→ purification by reduction with SO2 and re-oxidation with Cl2 or by adsorption and desorption on anion exchangers

Properties:

- solid grey-black crystals with metallic brilliance and a high tendency to sublimate (melting point: 113.6 °C, boiling point: 185.2 °C)

- molten iodine conducts electricity - rather low solubility in water (0.0013 mol /l = 33.88 g/l),

high solubility in iodide solutions and in organic solvents - toxic in elementary form, but essentially as ionic iodide - rather low reactivity, heavier reactions especially with P, Al, Fe and Hg, less

tendency to react with hydrogen - can be used as an oxidation agent and a reduction agent

Application:

- catalysts and very pure metals (van Arkel process for Zr and Ti via tetraiodides, used in the stereospecific polymerisation of butadiene)

- disinfections - pharmaceutical industry, food and feedstuff additives, agriculture - iodine impregnated activated carbon for Hg adsorption from waste gases - photography and rain cloud formation (AgI) - polyamide 6.6 (nylon) stabilisation

Properties of Halogen Compounds

In compounds fluorine has an oxidation number of –1 every time. Chlorine, bromine and iodine can reach oxidation numbers from –1 to + 7, whereby the electropositive character increases with the period number. Metal halogenides

- formation directly from the elements (partially very heavy reactions, even Au and Pt are attacked)

- Solubility in water and other polar solvents: high for salts formed with elements of the Ia and IIa groups low for salts formed with the heavy transition metals and the noble metals

- fluoride salts have partially inversed solubility properties compared to the other halogens

- neutral salts F-, “acid” salts [F-H-F]- (MeF · HF adducts)

Covalent halogen compounds with non-metallic elements

- formation from reaction of hydrocarbon compounds with fluorine (partially very heavy to explosive reactions) or by substitution reactions

- highest coordination numbers of positive “core” atom for fluorine compounds, e.g. SF6, PF6

- - solubility in water increases with the ionic character of bonds in the molecule - high volatility, low boiling temperatures especially in case of highly halogenated

compounds

Hydrogen Compounds of the Halogens

Formula HF

HCl HBr HI

Name of the poor compound

Hydrogen fluoride

Hydrogen chloride

Hydrogen bromide

Hydrogen iodide

Formation enthalpy ∆BH0 [kJ/mol]

-271 -92 -36 +26

ε0 (2 X-/X2) [V] (pH=0/14)

+3.05/+2.87 +1.63/+0.42 +1.06/+1.06 +0.54/+0.54

Reducing power

Oxidation power


colourless gas with

sticking smell, toxic

colourless gas with

sticking smell, toxic

colourless gas with

sticking smell toxic

colourless furning liquid

with sticking smell

toxic melting point [°C]

-83 -114 -87 -35

boiling point [°C]

+20 -85 -67 +26

Name of the aqueous solution

hydrofluoric acid

hydrochloric acid

hydrobromic acid

hydroiodic acid

pKs +3.2 -6.1 -8.9 -9.3 solubility [l/l H2O]

unlimited 507 612 425

Manufacturing and Use of Hydrogen Fluoride

Manufacturing in industry: Conversion of fluorspar CaF2 (Bayer Process)

CaF2 + H2SO4 → CaSO4 + 2 HF (200-250 °C)

Manufacturing in the laboratory:

Heating of acid fluorides of type MF ⋅ HF (e.g. M = K) MF ⋅ HF → MF + HF

Properties:

- highest bonding energy of all hydrogen compounds - hygroscopic liquid (melting point: - 83.36 °C,

boiling point: 19.51 °C) - soluble in water forming hydrogen fluoric acid (H3O

+F- - pKs = 3.2) - occurs in gas phase as (HF)6, at temperatures > 90 °C as HF - forms neutral salts MFx and acid salts MFx · (HF)n

Application:

- manufacture of inorganic fluorides (AlF3, BF3, UF6, NH4F) - manufacture of organic fluorocompounds

(esp. fluorohalogenhydrocarbons) - etching and polishing in the glass industry - manufacture of semiconductors

NOTE: Hydrogen fluoride and hydrogen fluoric acid attack glass and

quartz (SiO2 + 4 HF → SiF4 (g) + 2 H2O)! Store them only in Pb, Pt or in paraffin, PE, PP or Teflon bottles!

Industrial Important Fluorides

Aluminium Fluoride (AlF 3) - manufacture: Lurgi Process

2 Al(OH)3 → Al2O3 + 3 H2O (300-400 °C) Al2O3 + 6 HF → 2 AlF3+ 3 H2O (400-600 °C) Chemie Linz AG process 2 Al(OH)3 + H2SiF6 → 2 AlF3+ SiO2 + 4 H2O (100 °C)

- use: flux in the aluminium industry Sodium Aluminum Hexafluoride (Cryolite Na3AlF6)

- manufacture: 6 NH4F + 3 NaOH + 2 Al(OH)3 → Na3AlF6 + 6 NH3 + 6 H2O

- use: electrolytic manufacture of aluminium

Alkali Fluorides (NaF, KHF2, NH4F · HF) - manufacture: NaOH + HF or H2SiF6 - use: NaF – water fluoridation

KHF2 – frosting agent in glass industry, synthesis of F2 NH4F – oil extraction

Hexafluorosilicates (M2SiF6) - manufacture: 2 MCl + H2SiF6 → M2SiF6 + 2 HF - use: wood protection

Na2SiF6 - water fluoridation Uranum Hexafluoride (UF6)

- manufacture: UO2 + 4 HF → UF4 + 2 H2O UF4 + F2 → UF6

- use: separation of 235U and 238U in nuclear technology

Sulfur Hexafluoride (SF6) - manufacture: S + 3 F2 → SF6 - use: protective gas in high voltage installations

Boron Trifluoride (BF3) and Tetrafluoroboron acid (HBF4)

- manufacture: (1) Na2B4O7 + 6 CaF2 + 7 SO3 → 4 BF3 + 6 CaSO4 + Na2SO4 (reaction is carried out in conc. H2SO4) (2) HBO3 + 3 HF → BF3 + 3 H2O HBO3 + 4 HF → HBF4 + 3 H2O (reactions are carried out in conc. H2SO4)

- use: Friedel-Crafts catalyst in organic chemistry (BF3) galvanic metal deposition, fluxes, flame retardants

Manufacturing and Use of Hydrogen Chloride

Manufacturing in industry: (1) byproduct of synthesis of organic and inorganic chemicals (main source – 90 % of world market) e.g.: - manufacturing of chlorohydrocarbons (radicalic substituation) - reaction between amines and phosgene forming isocyanates R-NH2 + COCl2 → R–N=C=O + 2 HCl - substitution of chlorine by fluorine in organic molecules R-Cl + HF → R-F + HCl - manufacturing of phosphoric acid and of its esters - manufacturing of high surface silica by flame hydrolysis (SiCl4, H2, O2) 2 H2 + O2 → H2O SiCl4 + 2 H2O → SiO2 + 2 HCl (2) direct formation from the elements in a flame of 2000 °C

(Daniell burner - 8 % of world market) H2 + Cl2 → 2 HCl (3) byproduct of NaHSO4 formation from NaCl and H2SO4

(Leblanc process/Hargreaves process - 1-2 % of world market) SO2 + H2O + 0.5 O2 → H2SO4 (pre-process) NaCl + H2SO4 → NaHSO4 + HCl NaHSO4 + NaCl → Na2SO4 + HCl

Manufacturing in the laboratory:

2 NaCl + H2SO4 → 2 Na2SO4 + HCl Properties:

- well soluble in water (20 mol/l), short chain alcohols and ethers - traded concentrated hydrochloric acid is 38 % HCl in H2O - high oxidation power (e.g. forming chlorides from the elements)

Application:

- synthesis of chlorine containing organic compounds (addition reactions) - neutralisation reactions - acid hydrolysis reactions - regeneration of ion exchangers - polar solvent - manufacturing of chlorine (electrolysis/modified Deacon process) and chlorine

dioxide → Amount of HCl exceeds demand.

Manufacturing and Use of Hydrogen Bromide and Iodide

HBr HI Manufacturing of hydrogen halogenide:

(1) (2)

from the elements H2 + Br2 → 2 HBr (350 °C, Pt catalyst) Byproduct of organic bromine substitution reactions

(1) (2)

from the elements H2 + I2 → 2 HI (500 °C, Pt catalyst)

hydrazine + iodine N2H4 + I2 → 4 HI + N2

Manufacturing of halogenides:

MOH + HBr → MBr + H2O MOH + HI → MI + H2O or directly from the elements

Industrial application:

NaBr, CaBr2, ZnBr2

LiBr

KBr

NH4Br

use in oil industry

drying agent for air

photography

TiI4

NaI, KI AgI

catalysts

pharmaceutical purposes

photography induction of rain

Interhalogen Compounds

Electronegative partner

(valency = -1) F Cl Br

Electropositive partner

Valency

Cl +1 +3 +5 +7

ClF ClF3 ClF5

-

Br +1 +3 +5 +7

BrF BrF3 BrF5

-

BrCl - - -

I +1 +3 +5 +7

IF IF3 IF5 IF7

ICl (ICl3)2

- -

IBr - - -

Structure:

AB3 AB5 AB7

Lewis acids

AB3 AB5 AB7

Lewis acids Properties:

- synthesis from the elements (variation of reactant ratios and reaction conditions)

- similar to elements A2 and B2 - high fluoridation and oxidation activity

(increase with number of fluorine atoms, Cl > Br > I with respect to central atom) - disproportionation reactions of “middle” compounds,

e.g. 5 IF3 → I2 + 3 IF5 - high toxicity

Application: - ClF, ClF3, BrF3 and IF5 are used as industrial fluoridation agents (tons per year, e.g.

UF6 manufacturing) - ClF3 adducts with ammonia and hydrazine as fuel for rockets

Halogen Oxides – Overview

Valency F Cl Br I

-1 OF2

Oxygendifluorid1, (F-O-O-F)

Dioxygendifluorid1

- - -

+1 - Cl2O (Br2O) - +2 - (ClO, Cl2O2) - - +3 - (Cl2O3) (Br2O3) I4O9 (+3 and +4)

ClO2 +4 - (Cl2O4)

- I2O4

+5 - - (Br2O5) I2O5 +6 - (ClO3), Cl2O6 - I2O6

(+5 and +7) +7 - Cl2O7 - I2O7

1 NOTE: Compounds should be named as oxygen fluorides, NOT as oxides! Grey Fields – technical importance, () – not stable under standard conditions Properties:

- metastable endothermic explosive compounds compounds (without I2O5) - ionic character increases Cl < Br < I oxides - high oxidation activity (increase Cl < Br < I) - “in situ” utilisation - disproportionation reactions of “middle” compounds

Application:

- leaching agents - purification agents (oxygendonators) - fireworks

Halogen Oxides – Synthesis and Use

General synthesis:

- formation from the elements under consumption of energy (electrical discharges at deep temperatures)

- extraction of a water molecule from the corresponding acids - dis- and com-proportionation reactions

Special synthesises: a) Dihalogenmonoxides X2O

F2O: - hydrolysis of fluorine in basic solutions 2 F2 + 2 OH- → 2 F- + OF2 + H2O Cl2O: - formed by 2 Cl2 + 3 HgO → HgCl2 · HgO + Cl2O - use for synthesis of hypochlorites and chlorine isocyanates - leaching agent for textiles and wood Br2O: - 2 Br2 + 3 HgO → HgBr2 · HgO + Br2O (δ < -60 °C)

b) Halogenmonoxides (XO)n

ClO: - product of photolytic oxidation of Cl atoms in higher layers of the atmosphere

- radicalic properties (one free electron) - destroys ozone layer (ClO → Cl + O, Cl + O3 → ClO + O2)

c) Higher halogen oxides ClO2: - ER process by Erco, SVP process by Hooker (both starting from sodium chlorate) 2 HClO3 + SO2 → 2 ClO2 + H2SO4 (in 3-5 mol/l H2SO4) alternatively: NaClO3 + HCl → Cl2 + ClO2 - Munich or Kesting process:

1. Electrolysis of NaCl without cell separation NaCl + 3 H2O → NaClO3 + 3 H2

2. Reaction of chlorate solution with HCl 2 NaClO3 + 2 HCl → 2 ClO2 + Cl2 + 2 H2O + NaCl

- 2 NaClO2 + Cl2 → 2 ClO2 + 2 NaCl - Transport as sodium chlorite or stabilised with pyridine - Use as a leaching agent for wood pulp (no chlorolignin formation) and disinfection’s agent for potable water (less chlorination degree than Cl2)

I2O5: - Thermal treatment of iodine acid (200 °C) 2 HIO3 → I2O5 + H2O

Halogen Acids and Their Salts – Overview

Nomenclature: HXO – hypohalogenous acid XO- - hypohalogenite HXO2 – halogenous acid XO2

- - halogenite HXO3 – halogenic acid XO3

- - halogenate HXO4 – perhalogenic acid XO4

- - perhalogenate Acids: → protons are bonded with an oxygen atom Increase of acid strength

Valency F Cl Br I -1 (HOF) - - - +1 - (HClO)

Ks = 2.9 * 108 (HBrO)

Ks = 2.1 * 10-8 (HIO)

Ks = 2.3 * 10-11 +2 - - - - +3 - (HClO2)

Ks = 1.1 * 102 (HBrO2) (HIO2)

+4 - - - - +5 - (HClO3)

Ks = 5.0* 102 (HBrO3) Ks ~ 1

HIO3

+6 - - - - +7 - HClO4

Ks = 1010 (HBrO4) (HIO4)

(H5IO6) (H7I3O14)

Grey fields – technical importance, () – only stable in aqueous dilution, H5IO6 – ortho-periodic acid, H7I3O14 tri-periodic acid Base Anions: Increase of basic strength

Valency F Cl Br I -1 (OF-) - - - +1 - ClO - (BrO-) (IO-) +3 - ClO2

- BrO2- (IO2

-) +5 - ClO3

- BrO3- IO3

- +7 - ClO4

- BrO4- IO4

- H5-nIO6

n- H7-nI3O14

n-

Hypohalogeneous acids (HOX) and Their Salts (XO-)

HOF: - formation at –40 °C: F2 + H2O → HOF + HF

- decomposition in the gas phase and in weak basic solutions: 2 HOF → 2 HF + O2 - decomposition in neutral and acid solutions: HOF + H2O → HF + H2O2

HOCl/OCl- : - formation reactions:

(1) in water: Cl2 + H2O HCl + HClO (K << 1) (2) 2 Cl2 + 3 HgO + H2O → HgCl2 · HgO + 2 HOCl (3) in basic solutions: Cl2 + 2 OH- → Cl- + OCl- + H2O (industrial manufacturing with NaOH in solution at 40 °C or with Ca(OH)2 for solid salt – “Perchloron” process) (4) Olin /ICI /Thann and Pennwalt processes: Ca(OH)2 + 2 NaOCl + Cl2 +11 H2O → Ca(OCl)2 · NaOCl · NaCl · 12 H2O + Ca(OCl)Cl → Ca(OCl)2 2 H2O + 2 NaCl + 10 H2O (5) PPG process: Ca(OH)2 + 2 HOCl→ Ca(OCl)2 +2 H2O (6) electrolysis of seawater or brines in diaphragmless cells (small industrial consumers)

- acid not stable in higher concentrations (→ in situ use) - stable salts: LiOCl, Ca(OCl)2, Sr(OCl)2, Ba(OCl)2, NaOCl - commercial use for bleaching, for disinfections (e.g. water in swimming pools), neutralisation of poison gases and hydrazine manufacture

- use as “chlorinated trisodium phosphate” ([Na3PO4 · 11 H2O]4 · NaOCl) as cleaning agent in households and industry, especially in the USA - high oxidation power of the acid by formation of intermediate atomar

oxygen (HClO → HCl + O), oxidation potential ε0 (HClO/Cl-) = + 1.49 V

- very weak acid (Ks = 2.9 * 10-8), hydrolysis of salts - decomposition (catalysed by light) (1) acid solution: 2 HClO (aq) 2 HCl (aq) +O2

(2) basic solution: 3 HClO → 2 HCl + HClO 3

HOBr/OBr- : - formation and decomposition reactions similar to HOCl - disproportionation in water 2 BrO- → Br- + BrO3

-

- only alkali salts are stable until 0 °C HOI/OI- : - formation reactions: (1) 2 I2 + 3 HgO + H2O → HgI2 · HgO + 2 HOI

(2) in basic solutions: I2 + 2 NaOH → NaI + NaOI + H2O - very poor stability of acid, poor stability of salts - disproportionation: 5 HIO → HIO3 + 2 I2 + 2 H2O

Halogeneous acids (HOX2) and Their Salts (XO2-)

HClO2/ClO2

-: - acid decomposes 5 HClO2 → 4 ClO2 + HCl + 2 H2O - salt formation: 2 ClO2 + 2 MOH → MClO2 + MClO3 + H2O 2 ClO2 + 2 MOH + H2O2 → 2 MClO2 + O2 + H2O

- salts are relatively stable - high oxidation power, partially explosions - only industrial importance of NaClO2 formed by 2 ClO2 + 2 NaOH + H2O2 (excess) → 2 NaClO2 + O2 + H2O, used for ClO2 manufacture for small-scale users BrO2

-: - exists only as salts - formation of salts: (1) disproportionation of BrO- 2 BrO- → Br- + BrO2

- (2) comproportionation (solid reaction in absence of water) Br- + BrO3

- → 2 BrO2-

- decomposition in acid solutions, forming bromine HIO2/IO2

-: - both very unstable, no chemistry is known

Halogenic acids (HOX3) and Their Salts (XO3-)

HClO3/ClO3

-: - formation: (1) 2 HClO + ClO - → ClO3- + 2 HCl

(disproportionation of HClO in acid solutions) (2) 3 Cl2 + 6 OH- → ClO3

- + 5 Cl- + 3 H2O (3) Electrolysis of NaCl without cell separation NaCl + 3 H2O → NaClO3 + 3 H2 (technical process) - acid is stable up to 40 % - very strong oxidation agent in acid solution e.g. ClO3

- + 5 X- + 6 H+ → XCl + 2 X2 + 3 H2O - less oxidation power of basic salt solutions (in contrast to solid salts)

- industrial manufacture of Na salts by reaction (3) and metathesis of NaClO 3 with KCl (→ NaCl + KClO 3) - commercial use of NaClO 3 :

mainly for ClO 2 manufacture (ER and SVP processes), for synthesis of other ClO x compounds, as oxidation agent in uranium extraction (for U(IV) → U(VI)) and as herbizide

- commercial use of KClO 3 : fireworks and matches HBrO3/BrO3

-: - formation: (1) 3 Br2 + 6 OH- → BrO3- + 5 Br- + 3 H2O

(industrial process) (2) Electrolysis of NaBr without cell separation NaBr + 3 H2O → NaBrO3 + 3 H2 (technical process) (3) Oxidation with chlorine Br- + 3 Cl2 + 6 OH- → BrO3

- + 6 Cl- +3 H2O - acid stable up to 50 %, than decomposes to Br2, O2 and H2O - high oxidation power of the acid and the salts - used for redox titrations (colourless → red-brown, in acids) BrO3

- + 5 Br- + 6 H+ → 3 Br2 + 3 H2O -industrial application in flour treatment and in hair-setting lotions HIO3/IO3

-: - formation: (1) electrochemical or chemical oxidation of I2 e.g. I2 + 6 H2O + 5 Cl2 → 2 HIO3 + 10 HCl (2) MClO3 + I2 → Cl2 + 2 MIO3 (in hot HNO3, M = Na, K) (3) 3 I2 + 6 OH- → IO3

- + 5 I- + 3 H2O (4) NaIO3 + H2SO4 HIO3 + NaHSO4 - (1) and (2) are the commercial routes

- high stability of the acid and the salts - high oxidation power of the acid, moderate oxidation power of the salts

Perhalogenic acids (HOX4) and Their Salts (XO4-)

HClO4/ClO4

-: - formation: (1) Heating of alkali chlorates 4 MClO3 → 3 MClO4 + 2 MCl (industrial process in case of M = Na, metathesis reactions with NaClO4 to form the other perchlorates) (2) anodic oxidation of chlorates in basic solutions (technical process) ClO3

- + H2O → ClO4- + 2 H++ 2 e -

(3) electrolysis of chlorine in perchloric acid at 0 °C (Merck process): Cl2 + 8 H2O → 2 ClO4

- + 16 H+ + 14 e- (4) NaClO4 + HCl → HClO4 +NaCl - stable even as poor substance, salts in general stable - less oxidation power than chlorites, especially in case of diluted acid - low solubility of K, Rb and Cs salts - acid is traded in concentrations of 60-62 % in H2O - used in fireworks and as oxidation agent in rocket fuels BrO4

-: - formation: Oxidation of bromates with fluorine BrO3

- + F2 + H2O → BrO4- + 2 HF

- acid stable up to 55 %, pure salts are stable > 150 °C - less oxidation power because of kinetic hindering - decomposition at high temperatures only to BrO3

- HIO4/IO4

-: - formation: (1) Oxidation of iodates with chlorine IO3

- + Cl2 + H2O → IO4- + 2 HCl

(2) Thermal disproportionation of iodates Ba(IO3)2 → Ba5(IO6)2 + 4 I2 + 9 O2 - existence of periodate acid HIO4, H5IO6 – ortho-periodic acid and H7I3O14 tri-periodic acid (in water solution only H5IO6)

- anions in solutions: H4IO6-, H3IO6

2-, H2IO63-, IO4

-, H2I2O104-

- anions in salts of HIO4, H3IO5 (meso-acid), H5IO6, H6I2O10, H4I2O9 (di-periodic acids), H7I3O14


(1) Arnold Frederik Holleman, Egon Wiberg, Lehrbuch der anorganischen Chemie,



(7) P.W. Atkins, Physikalische Chemie, 2nd reprint of 1st edition, VCH Verlagsgesellschaft Weinheim, 1990

(8) Werner Büchner, Reinhard Schliebs, Gerhard Winter, Karl Heinz Büschel, Industrial Inorganic Chemistry, VCH Verlagsgesellschaft Weinheim, 1989

Pseudo Halogenes

Atom Groups

Atom group Hydrogen compound

Anion

CN N3 OCN SCN

cyanic acid – HCN nitrogen hydrogen acid - HN3 cyan acid - HNCO isocyan acid - HOCN thiocyan acid – HSCN isothicyan acid - HNCS

cyanide – CN- azide – N3

- cyanate – NCO- fulminate – OCN- thiocyanate – NCS- isothiocyanate – SCN-

Similar properties like halogens (Cl, Br. I) with respect to - acid reaction of hydrogen compounds

solubility in water (high solubility of alkali salts, low solubility of silver, mercury and lead salts)

- “oxidation number” = -1 - formation of singe bounded ligands in complexes - dimerisation to molecules X2 reacting like halogen molecules

(e.g. (CN)2, (NCS)2) - formation of interhalogen and inter- pseudohalogen compounds

(e.g. (NCS)Cl3, (NC)(NCS)) - com- and disproportionation reactions

(e.g. (CN)2 + 2 OH- → CN- + OCN- + H2O)


Chapter 6

Chalkogens (Oxygen Group)


Group Members

Oxygen (O)

Sulphur (S)

Selenium (Se)

Tellurium (Te)

Polonium (Po)


15.999 32.06 78.96 127.60 [209.98]

Discovery 1772 Scheele,

1774 Pristley

discoverer unknown (known since

antiquity)

1818 Berzelius

1782 von

Reichen-stein

1898 M. Curie


48.9 0.030 5 * 10-6 1 * 10-6 2 * 10-14 (radio-active)

melting point [°C]

-218.75 119.6 220.5

449.5 254

boiling point [°C]

-182.97 444.6 684.8 1390 962


colourless, odourless, tasteless gas (O2)

yellow non-metallic solid (S8)

red non-metallic (Se8) and grey metallic (Se8 ) solids

silver metallic solid

silver metallic solid

Electron negativity

3.5 2.5 2.4 2.1 2.0


-2 -2…+6 -2…+6 -2…+6 -2…+6

Reducing/ Oxidation Power

oxidising power

reducing

power

Metallic/ Non-metallic character

non-metallic

metallic

Acid/Basic properties of oxides

acid

base

Stability of valence states -2/-1 +2 +4 +6

Electron configuration: s²p4(d10) – need of accepting two electrons

or loosing 6 electrons for full saturation of electron shells

General Properties of Chalkogenes (1)

Hydrogen compounds and metal salts (-ides) - stability of hydrogen compounds decreases

H2O > H2S > H2Se >H2Te > H2Po

- bonding energy H-X decreases (463 kJ/mol (H2O), 348 kJ/mol (H2S), 276 kJ/mol (H2Se), 239 kJ/mol (H2Te))

- acid strength increases (pKs(25 °C = 15.74 (H2O), 6.92 (H2S), 3.77 (H2Se), 2.64 (H2Te))

- anions X2- are stable in case of all VIa group elements

- formation of “hydrogen per-compounds ” (H-X-X-H), “per-anions” (X22-) and

“polyanions” (Xn2-)

Halogen compounds (chalkogen halogenides)

- oxygen halogen compounds: O2F, O2F2, Hal2O, Hal2O3, HalO2, Hal2O5, HalO3, Hal2O7 (Hal = Cl, Br, I)

- sulphur and higher elements: formed in the compositions XnHal2 (n = 2, 3, 4), XHal2, XHal4 and XHal6 with X as

the electropositive partner - compounds can be formed from the elements, by com- and disproportionation

reactions and by treatment of oxides with halogenating agents (H-Hal, M-Hal)

General Properties of Chalkogenes (2)

Oxides

XO

X2O3 XO2

X2O5 XO3

oxidation state

+2 +3 +4

+5 +6

S SO2 – colourless gas

SO3 –colourless liquid

Se (SeO2)n – white needles, oxidising agent

Se2O5 SeO3 – colourless solid

Te TeO2 – yellow solid

Te2O5 TeO3 – yellow solid

Po PoO – black solid

PoO2 – yellow-red solid

PoO3 – only observed in traces

Acids/Bases

+2: +4: H2XO3 +6: H2XO4 H2SO3 –

pKs1/2 = 1.81/6.99, moderate reducing power

H2SO4 – pKs1/2 = -3/1.89, moderate oxidation power

H2SeO3 – strong acid pKs1/2 = 2.62/8.3

H2SeO4 – strong acid pKs2 = 1.74

- H2TeO3 – amphoteric pKs1/2 = 2.48/7.7, pKB1 = 2.7, low stability

o-H6TeO6 – pKs1/2 = 7.7/10.95, high oxidation power

Po(OH)2 - basic H2PoO3 - amphoteric -

oxidation power

acid strenght

Oxygen

Natural sources

- elementary: main component of air (20.5 %) - in compounds: water (88 %),

oxides and oxygen containing salts (e.g. SO42-, CO3

2-), essential part of biosphere Manufacturing

- in industry: air rectification (Linde process), electrolysis of water

- in laboratory: thermal or catalytic decomposition of peroxides Oxygen species

- neutral molecules: O2, O3 (ozone) - negative charged species: O2- (oxides - colourless),

O22- (peroxides - colourless),

O2- (hyperoxides - yellow),

O3- (ozonides - red)

- positive charged species: O2+ (dioxygenyl) in O2PtF6

Properties of O2

- colourless, tasteless and odourless gas - low solubility in water (3.05 l /100 l H2O) - essential for life (in dilution, toxic after long time exposure in elementary form) - high reactivity with near all elements, but only at high temperature or after catalytic

or photochemical activation, mostly strong exothermic reactions - reactivity is increased by adding small amounts of humidity

Application of O2

- generation of high temperatures (metallurgy, welding) - coal gasification - TiO2 production from TiCl4 - medicine - fuel cells

water cooling

O2

O3

water cooling

O2

O3

Ozone

Natural sources

- traces in atmosphere Manufacturing

- 3 O2 + 285.6 kJ/mol → 2 O3 - in industry/laboratory: activation of oxygen with

- thermal energy (>3500 K, very poor yields), - electrical energy

(“dark” discharges – ozonisator by SIEMENS), - photochemical energy (λ < 242 nm) or - chemical energy (e.g. oxidation of white phosphor)

- electrolysis of water, H2O2, HMnO4 - F2 + H2O → 2 HF + O, O + O2 → O3

ozonisator by SIEMENS Properties

- blue gas with characteristic odour - melting point: -192.5 °C, boiling point: -110.5 °C - well soluble in water (49.4 l/100 l H2O) - high endothermic, meta-stable compound - high oxidation power (O3 → O2 + O)

Application

- air disinfections - water disinfections - sterilisation of food

Ozone in the Troposphere (“Bad Ozone”)

Troposphere = lowest atmospheric layer up to 10 km height

Natural equilibrium between nitrogen oxides and ozone

“Photochemical smog” = anthropogenic increase - of NOx concentration - of hydrocarbon and CO concentration

→ Hydrocarbons, oxygenates and CO form peroxo radicals, which oxidise NO to NO2 instead of O3 (reaction 3)

emissions

emissionsUV radiation

hydrocarbons, CO

emissions

emissionsUV radiation

hydrocarbons, CO

Ozone in the Stratosphere (“Good Ozone”)

Stratosphere = atmospheric layer between 10 and 50 km height

Chapman cycle

- ozone formation: O2 + hν (< 242 nm) → Ο + Ο O + O2 + M (inert molecule) → Ο3

- ozone decomposition: O3 + hν (< 1200 nm) → Ο2 + Ο O3 + O → 2 Ο2

Catalytic ozone decomposition

O3 + X → O2 + OX OX+ O → X + O2

O3 + O →X 2 Ο2

(X = NO, H, OH – natural, Cl, Br – anthropogenic)

Peroxides

Industrial production of H2O2 - water dehydrogenation: electrolysis of sulphuric acid

2 H2SO4 → H2S2O8 + H2 (electrolysis), H2S2O8 +2 H2O → 2 H2O2 + 2 H2SO4

- oxygen hydrogenation: (I) 2 step antrachinone process (BASF)

(II) 1 step isopropanol process (Shell) isopropanol + O2 → H2O2 +acetone

Properties of H2O2

- meta-stable compound - decomposition: 2 H2O2 → 2 H2O + O2 + 196.2 kJ/mol

catalysed by noble metals, MnO2, high surface area particles, inhibited by acids as H3PO4 and organic acids

DO NOT STORE H2O2 IN GLASS BOTTLES !!!

- wide application as a clean oxidation agent - reducing properties with strong oxidation agents (e.g. Ag2O → Ag)

Technical application of H2O2

- leaching agent - production of perborates for detergents: NaBO2 + H2O2 → NaBO3 + H2O

Alkali metal peroxides

- Na2O2: production by 2 step oxidation of Na, strong oxidation agent, use in paper and textile leaching, use in respirators for CO2 removal Na2O2 + CO2 → Na2CO3 + 0.5 O2

- BaO2: production by thermal oxidation of BaO at 500-600 °C 2 BaO + O2 → 2 BaO2, use as igniting agent

The Ozone Leak in Antarctica (1)

1957 – begin of ozone measurements in Halley Bay

0

50

100

150

200

250

300

1960-1970

1984 1985 1986 1987

1974 – “Montreal protocol” = end of the use of fully halogenated hydrocarbons 1985 – discovery of the “ozone leak” 2002 – Stop of increasing ozone leak

Photos: “Magdeburger Volksstimme, Oct 12th 2002

The Ozone Leak in Antarctica (2)

(1) Formation of reservoir substances

- formation occurs in warmer areas - migration of the precursors to Antarctica

(2) Antarctic winter (-75…-85 °C)

- decomposition of reservoir substances in polar stratospheric clouds (PSC) by

catalytic reaction with ice and HNO3 · 3 H2O crystals

- formation of active chlorine

(3) Begin of the sunshine period (end of September)

(4) Antarctic spring

- warming of the atmosphere, changing of air pressure - mixing and replacing of Antarctic stratospheric air - chlorine content decreases, relaxation of the ozone layer

Sulphur (1)

Natural sources - in elementary form in sediments

(Italy, Poland, USA, Mexico, Peru, Chile, Japan) - in reduced form in sulphidic ores

(FeS2 – pyrit, CuFeS2, FeAsS, PbS, Cu2S, MoS2, ZnS, HgS, AsSx) - in oxidised form

(CaSO4 · 2 H2O – gypsum, CaSO4 – anhydrite, MgSO4 · 7 H2O, MgSO4 · H2O, BaSO4, SrSO4, Na2SO4 · 10 H2O)

Industrial production of elementary sulphur

- Mining - Extraction with superheated water and air under high pressure

(Frasch process)

sulphurpressured

air steam

sulphur containing limestone

steam

molten sulphur

rock

sulphurpressured

air steam

sulphur containing limestone

steam

molten sulphur

rock

- calcination of pyrite at 1200 °C under absence of air

(Outokumpu process) 83 kJ/mol + FeS2 → FeS + S

- Claus process use of H2S from desulphuration of natural gas, petrol, oil, synthesis gas or coke oven gas, 2 step process (1) H2S + 1.5 O2 → SO2 + H2O (non-catalytic combustion) (2) 2 H2S + SO2 → 3 S + 2 H2O (220-300 °C, alumina supported CoMo oxide catalyst, reactor cascade)

- COPE process = modified Claus process with partial reaction gas recycle Application of sulphur (50 * 106 t/a)

- Production of sulphur oxides/sulphuric acid (85-90 %), CS2 and P2S5 - vulcanisation of rubber - pharmaceuticals, exterminators - concrete and road building, paints - gunpowder and fireworks

Sulphur (2)

Sulphur modifications

meltingpoint

boilingpoint

solid,light

yellow

solid,near

colourless

119 °C – 159 °C liquid,

light yellow, low viscosity

159 °C - 243 °C liquid,

dark red-brown, high viscosity

243 °C – 445 °Cliquid,

dark red-brown, low viscosity

Chain length in solid and liquid state:

α-S and β-S S8 molecules

λ-S S8 moleculesπ-S Sn molecules (n = 5…30)µ-S polymerised molecules

gaseousred blue violet

meltingpoint

boilingpoint

solid,light

yellow

solid,near

colourless

119 °C – 159 °C liquid,

light yellow, low viscosity

159 °C - 243 °C liquid,

dark red-brown, high viscosity

243 °C – 445 °Cliquid,

dark red-brown, low viscosity

Chain length in solid and liquid state:

α-S and β-S S8 molecules

λ-S S8 moleculesπ-S Sn molecules (n = 5…30)µ-S polymerised molecules

gaseousred blue violet

Chemical properties

- reacts exothermally with most elements (without Au, Pt, Ir, N2, Te, I2 and noble gases) at moderate temperatures

- higher reactivity than oxygen - reacts with oxidising acids (to H2SO4) and alkaline solutions (forming

polysulphides - Sn2- - and thiosulphates – S2O3

2-) - inert in non-oxidising acids and in water - oxidation number of –2 in sulphides (S2-) formed with electropositive elements - oxidation number of +2, +4 and +6 in compounds with electronegative elements

(oxygen, halogens)

Sulphides

Hydrogen sulphide H2S

- natural sources: occurs in crude oil and natural gas, emitted from volcanos and mineral springs, biological decomposition of sulphur containing organic compounds

- synthesis: a) from the elements – H2 + S → H2S + 20.6 kJ/mol at 600 °C, MoS2 or alumina supported Co/MoOx catalysts (industrial process) b) treating of sulphides (e.g. pyrite) with hydrochloric acid FeS + 2 HCl → FeCl2 + H2S (laboratory method, using Kipp’s apparatus) c) obtained from purification of crude oil, natural gas and synthesis gas

- properties: stinking, colourless, high toxic gas, soluble in water, melting point: -85.6 °C, boiling point: -60.3 °C weak acid – H2S H+ + HS- 2 H+ + S2-,

moderate reducing agent - existence of hydrogen polysulphides H2Sn, and their metal salts

Industrial important metal sulphides

- Na2S, NaHS: prodced from Na2SO4 + C or from sodium polysulphide + Na amalgam, synthesis of organic sulphur compounds, depilatory in leather industry, ore flotation, precipitation of heavy metal ions

- K2S, NH4HS

Sulphur Oxides (1)

- existence of mono-sulphur oxides (SO, SO2, SO3 and SO4 – oxidation number +6) and poly-sulphur oxides (SnO, SnO2, S2O2, S3O9, (SO3-4)n)

- industrial relevance of SO2 and SO3 - SO2 and SO3 are anthropogenic emitted precursors of “acid rain”

Sulphur dioxide SO2

- Synthesis: (I) single or two stage combustion of elementary sulphur in air or pure O2 S + O2 → SO2 + 297 kJ/mol (II) calcination of sulphide ores (e.g. pyrite) in air or O2 using multiple hearth reactors, rotary kilns or fluidised bed reactors (650-1100 °C) 2 FeS2 + 5.5 O2 → Fe2O3 + 4 SO2 + 1655 kJ/mol, additional process step for removal of dust and catalyst poisons with respect to SO2 to SO3 oxidation (III) purification and evaporation of diluted waste acids (e.g. Venturi reconcentartion process, submerged- burner process, Pauling-Plinke process, Bayer-Bertrams process), yielding to 96 % acid (IV) SO2 extraction from wastes of exhaust air cleaning a) Müller-Kühne process (producing of cement) 4 CaSO4 + 4 SiO2 +2 C → 4 CaO · SiO2+ 4 SO2 + 2 CO2 (1400 °C) b) re-use of FeSO4 wastes from TiO2 manufacture (Bayer) 601 kJ/mol + FeSO4 → Fe2O3 + 2 SO2 + 0.5 O2 + H2O (900 °C)

- Properties: sticking odorous, colourless toxic gas, soluble in water (weak acid reaction), melting point: -75.5 °C, boiling point: -10.0 °C, reducing agent (forming SO3), main component of acid smog in winters

- Use: Production of sulphuric acid and sulphur containing chemicals (e.g. sulphites, dithionited, thiosulphates), disinfections agent (beer and wine industry), leaching agent

Sulphur Oxides (2)

Sulphur trioxide SO3

- Synthesis: (I) three stage catalytic oxidation of SO2 with air (contact process) 2 SO2 + O2 2 SO3 + 99 kJ/mol at 410-440 °C, catalyst: kieselgur supported V2O5

(V2O5 + SO2 → V2O4 + SO3, V2O4 + 0.5 O2 → V2O5) multi-stage fixed bed reactors

SO3 formation

SO3

yiel

d

SO3 decomposition

temperature [°C]catalyst

1st heat exchanger

1st tray(60 % conversion)

2nd tray(90 % conversion)

3rd tray(95 % conversion)

4th tray(98 % conversion)

catalyst

catalyst

catalyst

2nd heat exchanger

3rd heat exchanger

1

SO3 formation

SO3

yiel

d

SO3 decomposition

temperature [°C]catalyst

1st heat exchanger

1st tray(60 % conversion)

2nd tray(90 % conversion)

3rd tray(95 % conversion)

4th tray(98 % conversion)

catalyst

catalyst

catalyst

2nd heat exchanger

3rd heat exchanger

1

final step: sequential adsorption of SO3 and water in conc. H2SO4

(II) Nitrous process (lead chamber or tower process) at 80 °C, use of gaseous NO2 as the catalyst (N2O3 + SO2 → 2 NO + SO3, 2 NO + 0.5 O2 → N2O3) advantages: operates with lean reactant gas (0.5-3 % SO2), low operation temperature disadvantage: low final H2SO4 concentration (78 %) (III) Sulphur recycling from SOx containing wastes - Properties: 3 modifications – α-SO3 (= (SO3)p), β-SO3 (= (SO3)n),

γ-SO3 (= (SO3)3) – p > n > 3 melting points: α-SO3 62.2 °C, β-SO3 32.5 °C, γ-SO3 16.9 °C (depolymerisation of α-SO3 and β-SO3 during melting) boiling point: γ-SO3 44.4 °C, colourless, soluble in water, forming H2SO4 (strong acid reaction), oxidising agent (forming SO2),

- Application: production of H2SO4 and other sulphur compounds, production of alkyl sulphates (detergents)

Acids of Sulphur Oxides

Types of sulphate anions

sulphoxylate sulphite sulphate peroxo sulphate

thiosulphate disulphite disulphate peroxo disulphatedithionite dithionate

sulphoxylate sulphite sulphate peroxo sulphate

thiosulphate disulphite disulphate peroxo disulphatedithionite dithionate

- acids stable at high concentrations: sulphuric acid H2SO4, , disulphuric acid H2S2O7, peroxo sulphuric acid H2SO5, peroxo disulphuric acid H2S2O8, thiosulphuric acid H2S2O3

- other acids are stable only in dilution or as salts - general synthesis routes

reduction: 2 SO2 + 2 e - → S2O42-, 2 SO3 + 2 e - → S2O6

2- condensation: 2 HSO3

- → S2O62- + H2O,

2 HSO4- → S2O7

2- + H2O oxidation: 2 SO3

2- → S2O62- + 2 e -, 2 SO4

2- → S2O82- + 2 e -

- economically important acids: sulphurous acid H2SO3, sulphuric acid H2SO4

Sulphurous and Sulphuric Acid

Sulphurous acid H2SO3 Manufacture: SO2 + H2O [H2SO3] H+ + HSO3

- (K < 10-9) Properties: acid is stable only in dilution, salts are stable, moderate acid (KS1 = 1.54 · 10-2, KS2 = 1.02 · 10-9) reducing agent (forming H2SO4/ SO4

2-), can be oxidised by strong reducing agents e.g. 6 H+ + 6 e- + SO2 → H2S + 2 H2O (Zn/HCl), 2 SO2 + 4 H+ + 4 e- → S + 2 H2O (Fe2+), 2 SO2 + 2 H+ + 4 e- → S2O3

2- + H2O (HCOO-, S), Sulphuric acid H2SO4 Manufacture: contact process (see SO3) carried out in industry as a process unit outgoing from S, H2S or sulphidic ores (I) Oxidation of the starting material to SO2 (II) Oxidation of SO2 to SO3 (III) Formation of sulphuric acid in conc. H2SO4 SO3 + H2SO4 → H2S2O7 H2S2O7 + H2O → 2 H2SO4 Properties: very strong, oxidising acid (KS1 = 103, KS2 = 1.3 · 10-2), high affinity to water, strong heat accumulation during dilution, melting point: 3.0 °C (98 % acid), boiling point: 279.6 °C (100 % acid), azeotrope with water (98/2) boiling at 338 °C,

strong etching and oxidation agent, oxidises under SO2 formation organic substances to elementary carbon (coke), metals (without Pt and Au) to salts, hydrogen compounds to elements (HI → I2, H2S → S) weak reducing agent (forming peroxo disulphuric acid) oleum = solution of SO3 in H2SO4

Application: one of the most important base chemicals, production of fertilisers, mineral acids, inorganic and organic sulphates (e.g. detergents), uses as a catalyst (water removal), as electrolyte in batteries, as drying agent

Salts from Acids of Sulphur Oxides

Sodium hydrogen sulphite NaHSO3 Manufacture: NaOH+ SO2 → NaHSO3 Application: bleaching agent Sodium disulphite Na2S2O5 Manufacture: reacting NaOH + SO2 in a saturated NaHSO3 solution Application: photographic industry, paper industry, textile industry, food industry, water treatment Sodium disulphite Na2SO3 Manufacture: reacting NaOH + SO2 in a saturated Na2SO3 solution Application: photographic industry, paper industry, textile industry, food industry, water treatment Calcium hydrogen sulphite Na2S2O5

Manufacture: reacting limestone + SO2 Application: production of sulphite cellulose

Sodium thiosulphate Na2S2O3 and ammonium thiosulphate (NH4)S2O3

Manufacture: (I) 2 NaOH + SO2 + S → Na2S2O3 + H2O; Na2SO3 + S → Na2S2O3 (50-100 °C) (II) 2 Na2S + Na2CO3 + 4 SO2 → 3 Na2S2O3 +CO2 (III) 2 NH3 + SO2 + H2O → (NH4)2SO3, (NH4)2SO3 + S → (NH4)2S2O3 (80-110 °C) Application: fixing salts in photography (formation of [Ag(S2O3)]- and [Ag(S2O3)2]3- complexes soluble in H2O), anti-chlorination agent in bleaching plants and paper industry (Cl2 → Cl-), flue gas desulphurisation

Sodium dithionate Na2S2O4

Manufacture: (I) Zinc dust process (40 °C) Zn + 2 SO2 → ZnS2O4, Zn2S2O4 +2 NaOH → Zn(OH)2 +Na2S2O4 (II) Formate process (HCOO)- + OH- +2 SO2 → S2O4

2- +CO2 + H2O (III) Amalgam process

(IV) Sodium tetrahydroborate process Application: reducing agent in textile dying and printing starting material for sodium hydroxymethansulphinate (HO-CH2-SO2Na) used in direct and discharged printing

Other Important Sulphur Containing Compounds (1)

Disulphur dichloride S2Cl2 Manufacture: 2 S + Cl2 → S2Cl2 at 240 °C, catalysts: FeCl3 or AlCl3 Application: starting material for SOCl2 production, reaction with polyols gives additives for high pressure lubricating oils, catalyst for chlorination of acetic acid, vulcanisation of rubber Sulphur dichloride SCl2 Manufacture: S2Cl2 + Cl2 → 2 SCl2 at low temperatures, catalyst: I2 Application: starting material for SOCl2 production, sulphidising and chlorination reactions Thionyl chloride SOCl2 Manufacture: (I) reaction of SO2 or SO3 with Cl2, SCl2 and S2Cl2 over an activated carbon catalyst

(II) SO2Cl2 + PCl3 → SOCl2 + POCl3

Application: chlorination agent in organic chemistry (producing of herbicides, pesticides, pharmaceuticals, dyes and pigments),

non-aqueous electrolyte in galvanic cells Sulphuryl chloride SO2Cl2 Manufacture: SO2 + Cl2 → SO2Cl2, catalyst: activated carbon Application: chlorination and sulphochlorination agent (producing of herbicides, pesticides, pharmaceuticals, dyes and pigments)

Other Important Sulphur Containing Compounds (2)

Chlorosulphonic acid HSO3Cl Manufacture: SO3 + HCl → HSO3Cl in HSO3Cl Application: mild sulphonating and chlorosulphonating agent in organic chemistry Fluorosulphonic acid HSO3F Manufacture: SO3 + HF → HSO3F in HSO3F Application: fluorination agent in inorganic and organic chemistry (synthesis of sulphofluorides and sulphonic acids), catalyst for alkylation and polymerisation reactions, polishing agent for lead crystal glass Carbon disulphide CS2 Manufacture: (I) C + S2 → CS2 at 720-750 °C (II) CH4 + 2 S2 → CS2 + 2 H2S at 650-750 °C

Application: viscose industry (rayon), cellophan production, synthesis of CCl4, production of vulcanisation accelerators, flotation agents, corrosion inhibitors, herbicides and pharmaceuticals

Selenium and Tellurium

Selenium Manufacturing: 1. Oxidation of anode sludge from Cu electrolysis Ag2Se + O2 + Na2CO3 → Na2SeO3 + 2 Ag + CO2 2. Acidification with H2SO4 (→ separation of SeO3

2- and non-soluble TeO2) 3. Reduction of with SO2 H2SeO3 + 2 SO2 + H2O → Se + 2 H2SO4

Properties: 2 modifications in solid state - non-metallic red selenium Se8 - semi-metallic grey selenium Se8

Application: electronics (e.g. rectifier and photo cells, photocopiers Tellurium Manufacturing: 1. Oxidation of anode sludge from Cu electrolysis Ag2Te + O2 + Na2CO3 → Na2TeO3 + 2 Ag + CO2 2. Acidification with H2SO4 (→ precipitation of TeO2) 3. Resolving of TeO2 in base solutions 4. Chemical reduction of with SO2 TeO3

2- + 2 SO2 + H2O → Te + 2 SO42-

or electrochemical reduction Properties: silver-white colour with metallic brilliance,

semiconductor Application: additive in alloys of steel, copper, lead and tin (increase of mechanical properties)

Water

- covers 71 % of earth’s surface - 97 % of water is located in the oceans - essential part of plants (until 95 %) and animals (human: > 50 %)

water molecule - protons possess positive charge - electrons possess negative charge dipole momentum

104.9°

Physical properties

- increasing volume during freezing (ρwater = 0.9999 g/cm³, ρice = 0.9168 g/cm³) - maximum density at 3.98 °C (1.0000 g/cm³) - strong intramolecular hydrogen bridging bonds (high melting and boiling temperature)

p-T diagram chlarathe structure of ice

liquid water

water vapour

ice

temperature

pres

sure

liquid water

water vapour

ice

temperature

pres

sure

Chemical properties

- solves primarily ionic salts (dissociation of salts and solvatisation of the ions) and polar organic compounds (methanol, ethanol)

- autoprotolysis reaction 2 H2O H3O+ + OH- (K = 10-14)

- high thermal stability, low reactivity - acts normally as an moderate oxidation agent, with fluorine and other strong

oxidation agents as a reducing agent


Chapter 7

Pnictogens (Nitrogen Group)


Group Members

Nitrogen (N)

Phosphorus (P)

Arsenic (As)

Antimony (Sb)

Bismuth (Bi)


14.007 30.97 74.92 121.75 208.98

Discovery 1772 Scheele

1669 Brand

ca. 1250 Magnus

1492 Valentin

known since antique


0.017 0.10 1.7 * 10-4

2 * 10-5 2 * 10-5

melting point [°C]

-209.99 44.25 (P4) -

630.7 271.3

boiling point [°C]

-195.82 280.5 (P4) 616 (subli-

mation)

1635 1580


colourless, odourless, tasteless gas

white non-metallic solid (P4), red non-metallic solid (P8 ), black semi-metallic solid (P8 ), purple non-metallic solid (P8 )

yellow non-metallic solid (As4), black non-metallic solid (As8 ), grey metallic solid (As8 ),

grey metallic solid (Sb8 )

silver metallic solid (Bi8 )

Electron negativity

3.0 2.1 2.0 1.9 1.9


-3…+5 -3…+5 -3…+5 -3…+5 +3 (+5)


oxidising power

reducing

power


non-metallic

metallic


acid

base

Stability of valence states –3 +3 +5

Electron configuration: s²p3(d10) – need of accepting 3 electrons or loosing

5 electrons for full saturation of electron shells

General Properties of Pnictogens (1)

Valency states:

- - 3 with electropositive elements (stability decreases with increasing period number, -3 is unknown for Bi)

- + 3 and + 5 (+5 is favoured at lower, + 3 at higher period numbers) Hydrogen compounds and metal salts (-ides)

- stability of hydrogen compounds decreases NH3 (∆BH = – 46 kJ/mol) > PH3 (- 5 kJ/mol) > AsH3 (+ 66 kJ/mol) > SbH3 (+ 145 kJ/mol) > BiH3 (+ 278 kJ/mol)

- bonding energy H-X decreases NH3 (391 kJ/mol) < PH3 (227 kJ/mol) > AsH3 (297 kJ/mol) > SbH3 (257 kJ/mol) > BiH3 (194 kJ/mol)

- basic strength decreases (NH3 (pKB(25 °C) = 4.75) < PH3 (27), AsH3, SbH3 and BiH3 are not stable in aqueous solutions)

- N and P form a wide variety of compounds with more than one pnictogen atom (e.g. N2H4, P3H5)

Halogen compounds (pnictogen halogenides)

- nitrogen: formed in the compositions NX3, N2X4, N2X2 and N3X as derivates from the hydrogen compound

- higher elements: XHal3, X2Hal4 and XHal5 with X as the electropositive partner - formation of oxyhalogenides (e.g. NOCl, POCl3)

General Properties of Pnictogens (2)

Oxygen compounds (pnictogen oxides)

- nitrogen: formed in the compositions N2O, NO, N2O2, N2O3, NO2, N2O4, N2O5, NO3 and N2O6

- higher elements: X2O3 and X2O5 with X as the electropositive partner Acids/Bases

- nitrogen: H3NO, HNO, H2N2O2, H2N2O3, HNO2, HNO3, NO4

3- and HNO4 - phosphorus: H3PO2, H4P2O4, H3PO3, H2P4O5, H2P4O6, H3PO4, H4P2O7 H4P2O8 and

H3PO5 - higher elements: H3XO3 and H3XO4 - acidity decreases with increasing period number, acids containing X(V) are stronger

than acids with X(III) - N, P, As and Sb oxides form acids, Bi2O3 possess basic properties

Sulphur compounds (pnictogen sulphides)

- nitrogen: sulphur nitrides (e.g. S4N4), sulphur nitride halogenides

(e.g. (NSX)n), sulphur nitride oxides (e.g. (NSXO)3) and sulphuric acid derivates of hydrogen compounds (e.g. NH2SO3H – amido sulphuric acid)

- phosphorus: X4Sn (P: n = 2-10, As: n = 3-10, similar composition for selenids) - higher elements: X2S3 and X2S5, for As, Sb and Bi with metal sulphide character

Nitrogen

Natural sources

- elementary: main component of air (79.5 %) - in compounds: water (solved nitrate salts – 1 % of total N),

salt deposits (Chile, India) essential part of biosphere (amino acids) -

0.001 % of total N Manufacturing

- in industry: air rectification (Linde process), - in laboratory: thermal or catalytic removal of oxygen from air

(product contains noble gases), pure nitrogen by oxidation of ammonia (NH3 + HNO2 → N2 + 2 H2O)

Nitrogen species

- N2, N3- (nitride), N3

- (azide) Properties

- colourless, tasteless and odourless gas - triple bond between the nitrogen atoms - low solubility in water (3.05 l /100 l H2O) - essential for life (formation of amino acids, assimilation for most plants as NH4

+, NO2

-, NO3- and urea, special bacteria – “azobacter” - can convert gaseous N2)

- very poor reactivity with near all elements, only at high temperature, mostly strong endothermic or kinetically hindered reactions

Application of N2

- inert purging gas - cooling agent (in liquid state) - synthesis of ammonia, hydrazine, hydoxylamine and nitric acid

Important Nitrogen Hydrogen Compounds (1)

Ammonia NH3

- Manufacturing

from the elements: N2 + 3 H2 2 NH3 - Properties

colourless toxic gas with sticking odour, melting point: -77.8 °C, boiling point: -33.4 °C, high solubility in water (702l/l H2O at 20 °C), forming weak basic solutions (NH3 + H2O NH4

+ + OH-, pKB = 4.75)

- Application production of fertilizers (80 %, incl. fertilizers from nitric acid), plastics (10 %), explosives (5 %), herbicides, organic chemicals one of the most important industrial chemicals

Hydrazine H2N–NH2

- Manufacturing

(1) oxidation of ammonia with hypochlorites (in-situ formation from Cl2): 2 NH3 + ClO- → N2H4 + H2O + Cl-

(2 step Raschig process in base solution, 2 step Bayer process in acetone) (2) oxidation of ammonia with H2O2 and methyl ethyl ketone (2 step Pechiney Ugine Kulmann process) (3) 2 step oxidation of urea with hypochlorites – presently not in commercial use

- Properties colourless, fuming liquid with high viscosity and strange odour, melting point: 2.0 °C, boiling point: 113.5 °C, forming a hydrate N2H4 · H2O (high viscose liquid with “fishlike” odour, melting point: - 51.7 °C, boiling point: 118.5 °C), endothermic meta-stable compound, decomposition only at high temperatures, soluble in water with basic reaction ((H3N-NH2)

+ +OH- - Application

synthesis of a large amount of organic chemicals, of polymerisation initiator, of herbicides and of pharmaceuticals; acts in water as a corrosion inhibitor

Important Nitrogen Hydrogen Compounds (2)

Hydroxylamine NH2OH - Manufacturing

(1) modified Raschig process (3 step process) HNO2 + 2 H2SO3 + H2O → NH2OH + 2 H2SO4, carried out with NH4NO2 and SO2 in diluted H2SO4 (2) NO reduction process (BASF, Iventa – favoured process) 2 NO + 3 H2 → 2 NH2OH (in H2SO4, catalyst: Pt or Pd on C) (3) Nitrate reduction process (3 step process using NH4NO3, hydrogen, phosphoric acid and cyclohexanone)

- Properties colourless, odourless solid (needles), melting point: 33 °C, decomposes at moderate temperatures to NH3, N2 and H2O, stable only in the absence of air or in aqueous solutions, weak basic properties (pKB = 8.2), salts (e.g. sulphates) posses much higher stability

- Application production of caprolactam (97 %) for synthetic textiles, of pharmaceuticals, herbicides and l acquers, anti-oxidising agent

Nitrogen hydrogen acid HN3

- Manufacturing (1) from amides (obtained by the reaction of alkali metals with ammonia: 2 Na + 2 NH3 → 2 NaNH2 + H2) and dinitrogen monoxide: NaNH2 + N2O → NaN3 + H2O (190 °C) (2) from amides and nitrates: 3 NaNH2 +NaNO3 → NaN3 + 3 NaOH + NH3 (100 °C, high pressure, in liquid NH3) (3) from salpeterous acid and hydrazine: HNO2 + N2H4 → HN3 + 2 H2O (0 °C, in ether) 90 % acid can be obtained by distillation with diluted H2SO4 and water removal with CaCl2

- Properties colourless, low viscose and high toxic liquid with piercing, unbearable odour, melting point: - 80 °C, boiling point: 35.7 °C, strong endothermic meta-stable compound, explosion at high temperatures and after blow (decomposition to the elements) soluble in water with weak acid reaction (H3O

+ +N3-, pKS = 4.92)

forms salts with properties similar to chlorides - Application

Pb(N3)2 - made from Pb(NO3)2 and NaN3 - is used in explosives and in air bags (cars).

Industrial Ammonia Synthesis 1. A Synthesis of High Complexity

Starting materials - air (nitrogen and oxygen) - water - natural gas, oil or coke

Process scheme of an integrated ammonia plant

air natural gas heavy oil water coal

rectification steam reforming

partial oxidation

gasification

nitrogen oxygen

rough synthesis gas (N2, O2, CO, H2, H2S)

absorption of CO and H2S

adsorption of CO2 and H2O on zeolithes

scrubbing with liquid N2 at – 196 °C and 80 bar

final synthesis gas (N2, H2)

ammonia synthesis

separation and purification of ammonia

final product (NH3)

air natural gas heavy oil water coal

rectification steam reforming

partial oxidation

gasification

nitrogen oxygen






ammonia synthesis


final product (NH3)






ammonia synthesis


final product (NH3)

Industrial Ammonia Synthesis 2. Pre-Processing Steps (1)

Steam reforming - raw materials: natural gas, naphta, water - 2 C2H2n+2 + n H2O → n CO + 2(n+1) H2 - Process steps:

(1) Desulphurisation of raw materials by hydrogenation over CoO or NiO/MoO3 catalysts at 350 – 450 °C (2) Adsorption of formed H2S on ZnO H2S + ZnO → ZnS + H2O (3) Primary reforming with steam at 700 – 830 °C and 40 bars over NiO/Al 2O3 catalysts (4) Secondary reforming of methane at 1000 – 1200 °C CH4 + H2O → CO + 3 H2 (5) Adjustment of stochiometric N/H ratio by feeding air into the second reformer

Partial Oxidation

- raw materials: heavy fuel oil, air (enriched with oxygen) - 2 C2H2n+2 + n O2 → 2n CO + 2(n+1) H2 - non-catalytical process at 1200 – 1500 °C and 30 – 40 bar - advantage: no desulphurisation step,

disadvantage: need of additional O2 (= additional air rectification step)

Coal gasification - raw materials: coal, air, water - C + H2O CO + H2, 2 C + O2 2 CO,

C + O2 → CO2 for heat generation (1/3 of all coal is oxidised totally) - 1200 °C/solid-bed reactor (Lurgi process) or

800 – 1100 °C/fluidised bed (Winckler process) or 1400 – 1600 °C and 1 bar/fly ash (Koppers Totzeck process)

Industrial Ammonia Synthesis 2. Pre-Processing Steps (2)

Conversion of carbon monoxides (~ 7 % in the feed) - CO poisons ammonia synthesis catalyst - CO + H2O → CO2 + H2 - 350 - 380 °C/FeCrOx catalyst (rest: 1 g CO /m³ air)

350 - 380 °C/sulphur insensitive Co/Mo oxide catalyst (rest: 0.3 % CO)

Removal of CO and H2S - absorption with organic solvents (Rectisol process) - absorption with K2CO3 (Benfield process) - combination of both process modifications

Final purification

- aim: remove of any oxygen containing compounds (CO, CO2, H2O, O2) and of H2S - adsorption of CO2 and H2O on zeolithes - scrubbing with liquid N2 at – 196 °C and 80 bar (condensing of hydrocarbons, enrichment

of N2 if necessarily)

Final composition of synthesis gas - H2: 74.0 %, N2: 24.7 % (H2/N2 = 3),

CH4: 1.0 % Ar: 0.3 %, CO + CO2 < 10 ppm

Catalyst composition and preparation - promoted α-iron catalyst - composition of starting mixture:

94.3 % Fe3O4, 0.8 % K2O, 2.3 % Al2O3, 1.7 % CaO, 0.5 % MgO, 0.4 % SiO 2

- preparation: (1) melting a mixture of magnetite (Fe3O4) and additives at 1600 – 2000 °C (2) rapid cooling (3) crushing and sieving (required particle diameter: 6-10 mm) (4) in-situ reduction in the reactor with synthesis gas at 350 – 400 °C and 70 – 300 bar (Fe3O4 + 4 H2 → 3 Fe + 4 H2O)

- Fe - active component - K2O acts as electronic promoter (increase of activity) - Al2O3 and SiO2 prevent sintering and provide acid/base sites - CaO increase resistance against S and Cl - other additives: oxides of Li, Be, V and U - lifetime > 10 years

Industrial Ammonia Synthesis 3. Synthesis and Purification of Ammonia

equilibrium limited, exothermic reaction with volume decrease

Equilibrium diagram

tubular reactor with integrated heat exchanger

multiple-bed reactor with integrated heat exchanger

temperature tubular reactor multiple-bed reactor

pressureresistent

outer tubecatalytic contactheat

exchanger tube

temperature tubular reactor multiple-bed reactor

pressureresistent

outer tubecatalytic contactheat

exchanger tube

Process conditions

process

country pressure [bar] temperature [°C]

Haber Bosch process

Casale process Fauser process Claude process

Mont Cenis process Kellog process

Germany

Italy Italy

France Germany

USA

200

600-800 200-300 900-1000

100 160-240

500 500 500

500-600 400-450

500

Product separation and purification

- obtained yield: 10-15 % - product separation by freezing out the ammonia (-25 °C)

and re-vaporisation or by absorption in water - recycle of the unconverted reactants after removal of water - compression and cooling for storage

Nitrogen Oxides and Acids – Overview

Oxdiation number of N +1 +2 +3 +4 +5 Oxides N2O NO,

N2O2 N2O3 NO2,

N2O4 N2O5

Acids HNO, H2N2O2

N2O32- HNO2 HNO3

NO43-

HNO4

Oxides:

- endothermic meta-stable compounds (without N2O4 and N2O5) - occur during high temperature combustion processes from nitrogen oxidation

(instead of N2O) - equilibriums NO/N2O2 and NO2/N2O4 (dimerisation primarily at higher

temperatures) - N2O, NO/N2O2 and NO2/N2O4 have a large technical importance and environmental

relevance as anthropogenic emissions (→ ozone and as precursors for “acid rain”)

Acids:

- acid strength increases with increasing number of oxygen atoms - HNO2 (only stable in gas phase or in aqueous dilution – salts are stable) and HNO3

are stable and of large technical importance - H2N2O2 decomposes at room temperature within days, salts are stable - HNO and HNO4 are meta-stable at low temperatures and decompose under normal

conditions - N2O3

2- and NO43- exist only as salts

Nitrogen Oxides

Dinitrogen monoxide N2O (NNO)

- Industrial manufacturing careful heating of ammonium nitrate or or of a mixture of NH3 and HNO3 or of a mixture of sodium nitrate and ammonium sulphate (forming NH4NO3) at 200 °C (danger of explosions!) NH4NO3 → N2O + 2 H2O + 124.1 kJ/mol,

product of biological nitrification and denitrification processes - Properties

colourless gas with weak sweet odour and intoxicating effect, melting point: - 90.9 °C, boiling point: - 88.5 °C, well soluble in water (0.60 l/l H2O) and fats,

oxidation agent, supports combustion processes similar to oxygen - Application

narcotic agent, propellant in ice-cream and whipped cream Nitrogen monoxide NO

- formed by nitrogen oxidation at temperatures > 1500 °C or by Pt catalysed short contact time combustion of ammonia (Ostwald process: 4 NH3 +5 O2 → 4 NO + 6 H2O)

- high endothermic compound (∆BH = 180.6 kJ/mol), molecule contains one unpaired electron = free radical (!)

- colourless, high toxic gas, low solubility in water - melting point: - 163.6 °C, boiling point: - 151.8 °C - dimerises especially in solid and liquid state to N2O2, - rapid oxidation to NO2 under presence of air

(< 650 °C – equilibrium limited reaction) - supports combustion processes providing its oxygen - forms nitrosyl compound with halogens (NO-X and NOF3) - NOx removal from waste gases by SCR with ammoinia or

by reaction with CO and hydrocarbons (3 way catalyst)

Nitrogen dioxide NO2 /dinitrogen tetroxide (O2N–NO2) - formed by NO oxidation at moderate temperatures

2 NO + O2 2 NO2 + 114.2 kJ/mol - in equilibrium with N2O4

2 NO2 (brown-red) N2O4 (colourless) + 57,2 kJ/mol - characteristic odour, high toxicity - melting point: - 11.2 °C (0.01 % NO2/99.99 N2O4),

boiling point: +21.2 °C (20 % NO2/80 N2O4) - oxidation agent stronger than N2O and O2, supports combustion - forms nitryl compounds (O2N-X) with halogens

Industrial manufacturing of Nitric Acid

OSTWALD Process (3 step process for 50 - 68 % acid) (1) catalytic combustion of NH3 over Pt-Rh alloy gauzes (→ short contact times) at 820 – 950 °C and 1 – 12 bar 4 NH3 + 5 O2 → 4 NO + 6 H2O + 904 kJ/mol (2) further oxidation of NO to NO2/N2O4 at ca. 150 °C 2 NO + O2 → 2 NO2 + 114 kJ/mol, 2 NO2 → N2O4 + 57 kJ/mol (3) absorption in water under presence of excess air at up to 15 bar: 2 NO2 (=N2O4) + H2O + 0.5 O2 → 2 HNO3

Pt-Rh alloy gauze

NO

air +ammonia

Pt-Rh alloy gauze

NO

air +ammonia

Direct strong nitric processes for highly concentrated acid (Europe)

- CAN process (Uhde) (1) catalytic combustion of NH3 at atmospheric pressure (2) oxidation of NO at 1.6 bar (3) physical absorption of NO in highly concentrated frozen HNO3 (4) rectification of rough acid to 98 – 99 % acid (sump) and N2O4 (head) (5) conversion of N2O4 from (4) with pure oxygen in diluted HNO3 to 98 – 99 % acid

- SABAR process (Davy McKee’s) (3) absorption of NO in azeotropic HNO3 (68 – 69 %) in the presence of oxygen at 6 – 13 bar (4) blowing out of the acid with secondary air, acid rectification (sump: azeotropic acid – recycled, head: near pure acid)

Indirect extractive distillation processes (USA)

- sulfuric acid process counter-current extractive distillation of rough HNO3 (e.g. from Ostwald process) with concentrated H2SO4

- magnesium nitrate process distillation of rough HNO3 (e.g. from Ostwald process) with a 72 % Mg(NO3)2 solution

Purification of waste gases

- alkali scrubbing with solutions of NH3, NaOH or urea - thermal (> 1000 °C - NSCR) or catalytic reductive post-combustion (170 –

600 °C - SCR) with reducing agents (e.g. hydrocarbons, hydrogen, CO, NH3)

Properties and Application of Nitric Acid

Properties - colourless liquid, decomposes slowly (faster at higher temperatures) to nitrous

oxides and water: 2 HNO3 → 2 NO2 + H2O + 0.5 O2 - melting point: - 41.6 °C, boiling point: + 82.6 °C - azeotrope with of 69.2 % HNO3 with water (boiling point: 121.8 °C) - azeotrope = traded “concentrated HNO3” - strong oxidation agent, reacts with Cu, Ag, Hg, S, P and organic substances (not

with Au, Pt, Rh, Ir) 4 H+ + NO3- + 3 e - NO + 2 H2O

- mixture of HNO3 and HCl (1 : 3) oxidises even oxidises Au (HNO3 + 3 HCl → NOCl + 2 Cl (atomic) + 2 H2O)

- mixture of HNO3 and H2SO4 (1 : 9) is used as a nitration agent in organic chemistry - strong acid (pKs = - 1.44) - salts (nitrates) have a high solubility in water, low melting points (250 – 350 °C)

and decompose easily in the heat: alkali and earth alkali metal nitrates: KNO3 → KNO2 + 0.5 O2, transition metal nitrates: Cu(NO3)2 → CuO + 2 NO2 + 0.5 O2,

ammonium nitrate: NH4NO3 → N2O + 2 H2O + 124.1 kJ/mol (200-260 °C) or NH4NO3 → N2 + 0.5 O2 + 2 H2O + 206.2 kJ/mol (>300 °C)

Application

- 75 - 85 % production of ammonium nitrate (NH4NO3) for fertilizers (80 %), explosives (~ 20 %) and N2O synthesis (from 50 – 70 % acid)

- 10 % for production of adipic acid (HOOC-(CH2)4-COOH - fiber and plastic precursor)

- 3 % production of TNT (with high concentrated acid) - 3 % nitration of benzene (aniline precursor, reaction is carried out

high concentrated acid) - 2 % alkali and earth alkali nitrates (fertilisers) - 1 % organic nitro-compounds

Elementary Phosphorus (1)

Natural sources - occurs only in compounds - 0.1 mass % of earth, 13th common element - inorganic sources: phosphates (apatite Ca3(PO4)2 · CaX2 , X = OH, F, Cl;

iron and aluminium phosphates) - biological importance: participation in metabolism processes as

phosphorus acid esters and phosphates (e.g. ADP/ATP)

Manufacturing in industry

- highly endothermic electro-thermal reduction of phosphates with coke and quartz at 1400 – 1500 °C: 1542 kJ/mol + Ca3(PO4)2 + 3 SiO2 + 5 C → 3 CaSiO3 (slag) + 5 CO + P2 (g)

(reduction of P2O5 by C, SiO2 is added to form a slag with Ca) - condensation and distillation of rough phosphorus

→ “white phosphorus”

carbon electrode

carbon electrodes

outletfor slags

outletfor slags

phosphate

phosphatite

electrode mass

outletfor moltenironfromelectrodemass

gas outlet (P2)carbon electrode

view from side view from top

carbon electrode

carbon electrodes

outletfor slags

outletfor slags

phosphate

phosphatite

electrode mass

outletfor moltenironfromelectrodemass

gas outlet (P2)carbon electrode

view from side view from top

Elementary Phosphorus (2)

Modifications

White Phosphorus

Red Phosphorus

Violet Phosphorus

Black Phosphorus

Manufacture electro-thermal reduction of phosphates

thermal treatment of white P

at 200-400 °C for 20-30 hours

(∆BH = - 17.7 kJ/mol)

thermal treatment of

white P at > 550 °C

for 1-2 weeks

thermal treatment of white P

at 380 °C for some days (adding of

dispersed Hg as catalyst)

Molecular and crystal structure

P4 tetrahedrons P8 , amorphous

P8 , complex layer

structure

P8 , layer structure

(similar to graphite)

Stability meta-stable at 25 °C,

thermodyn. stable at > 620 °C

meta-stable at 25 °C

meta-stable at 25 °C,

thermodyn. stable

at 550-620 °C

stable form at < 550 °C

Melting point 44.3 °C sublimation at ~ 580 °C

sublimation at 620 °C

changes to violet P

Boiling point 280.5 °C - - - Toxicity high (T+) low low low

Reactivity extremely high, high flammable, strong reducing

agent

moderate, heavy reaction

only with strong oxidation agents

low, flammable at

> 400 °C

low, flammable at

> 400 °C

Metallic/ non-metallic

character, electrical

conductivity

non-metallic isolator



metallic semiconductor

Phosphorous species in compounds

- occur in all oxidation states from –3 to +5

Application - 90 % for manufacture of P2O5 (→ phosphorous acid, phosphates) - synthesis of P-S and P-Halogen compounds (→ organic chemistry) - safety matches (red phosphorous) - military purposes

Phosphorus Oxides

Phosphorus trioxide (P4O6) and phosphorus pentoxide (P4O10)

phosphorus trioxide (P4O6) phosphorus pentoxide (P4O10)

- manufacture: combustion of P4 at low temperatures and oxygen

shortage (P4O6) or with dried air under O2 excess (P4O10) P4 + 3 O2 → P4O6 + 1641.2 kJ/mol P4 + 5 O2 → P4O10 + 2786 kJ/mol, separation of oxides by fractioned vaporisation

- properties: P4O6 – white wax-like high toxic solid, melting point: 23.8 °C, boiling point: 175.3 °C (under N2), oxidation to P4O10 at temperatures > 70 °C (∆RH = + 672.4 kJ/mol), reaction with cold water to phosphonic acid (H3PO3), with HCl to H3PO3 and PCl3 and with halogens to phosphoryl halogenides P4O10 – white snow-like high toxic solid, sublimation at 358.9 °C, strong hygroscopic, reacts with water to phosphorus acid (H3PO4), very weak oxidation agent (only at high temperatures)

- application: P4O6 has no technical importance, P4O10 is used as a drying and dehydrogenation agent and (mainly) for the manufacture of phosphorus acid and its esters.

Phosphorus Acids

Formal “Oxidation

number” of P “+1” “+2” “+3” “+4” +5

acids

H3PO2 H4P2O2 (HPO2)n, H3PO3, H4P2O5

- H4P2O6 (HPO3)n, H3PO4, H3PO5, H4P2O7, H4P2O8

phosphinate phosphonate phosphate peroxophosphate diperoxophosphatephosphinate phosphonate phosphatephosphinate phosphonate phosphate peroxophosphate diperoxophosphateperoxophosphate diperoxophosphate

anions of ortho-phosphorus acids (P=O double bond is delocalised between all P-O baonds)

Deprotonation in water

phosphinic acid (=hypophosphoric acid) = one-base acid phosphonic acid (= phosphorous/phosphoric acid) = two-base acid

phosphorus acid, peroxophosphorus acid and diperoxophosphorus acid = three-base acids

hypodiphosphonate(diphosphate(III))

diphosphonate(diphosphate(III))

hypodiphosphate(diphosphate(IV))

diphosphate(diphosphate(V))

peroxo-diphosphate(V)

diphosphate(III,V)diphosphate(II,IV)













anions of diphosphorus acids

Technical important phosphorus acids: H3PO4, H3PO3 and H3PO2

Phosphorus acid - H3PO4

Manufacturing (I) “Wet processes” Dihydrate process (80 °C → CaSO4 · 2 H2O),

hemihydrate process (95 °C →) CaSO4 · ½ H2O) Ca3(PO4)2 (apatite) + H2SO4

→ 3 CaSO4 + 2 H3PO4, concentration of the acid by vacuum evaporation or submerged burners,

purification of the acid by precipitation and extraction with organic solvents

(II) “Thermal process” Oxidation of white P4 in air excess and absorption of the formed P4O10 together

with water in conc. H3PO4 (85 %) P4 + 5 O2 → P4O10, P4O10 + 6 H2O → 4 H3PO4

(single tower IG process, double tower TVA process)

Properties pure acid = colourless, clear, odourless hard solid,

melting point: 42.4 °C, partially condensation (at 200 °C completely)

2 H3PO4 → H4P2O7 + H2O concentrated (85 %) acid = high viscous liquid, melting point: 21.1 °C, boiling point: 158 °C,

three-base middle-strong acid (pKS1 = 2.16, pKS2 = 7.21, pKS3 = 12.32),

pure acid = strong oxidation agent > 400 °C, diluted acid = no oxidising properties

pH

buffer area

pH

buffer area

Application producing of salts (phosphates of K, Na and NH4), use of acid itself in cleaning agents, in metal treatment and

polishing, as an acidification agent in soft drinks (colas, lemonades),

organic chemistry (e.g. esterification)

H3PO3 and H3PO2

Ortho-phosphorous/ortho-phosphoric/ortho-phosphonic acid (H3PO3) - manufacture: (I) PCl3 + 3 H2O → H3PO3 + 3 HCl at 190 °C

(II) P4O6 + 6 H2O → 4 H3PO3 - properties: colourless crystals, melting point: 73.8 °C,

middle-strong two base acid (pKS1 = 2.0, pKS2 = 6.6), in aqueous solutions high solubility of alkali salts, low solubility of other salts, hydrogen phosphonates dimerise in the heat 2 H2PO3

- → (HO2–P–O–PO2H)2- + H2O, strong reducing agent, reduce noble metals from

their salts, halogens to halogenides, H2SO4 to H2SO3, stable under air atmosphere at room temperature

disproportionation at heating of the dry acid 4 H3PO3 → PH3 + 3 H3PO4 (130 – 140 °C)

- application: reducing agent, industrial synthesises of base lead phosphonate (PVC stabilisator), of organic phosphonic acids and phosphorous acid esters

Ortho-phosphinic/ortho-hypophosphoric acid (H3PO2)

- manufacture: (I) Cooking of white P4 with NaOH or Ca(OH)2 2 P4+ 3 Ca(OH)2 + 6 H2O → 2 PH3 + 3 Ca(H2PO2)2↓ Ca(H2PO2)2 + H2SO4 → CaSO4↓ + H3PO2, isolation by vaporising of the solution or by extraction with diethyl ether

(II) PH3 + 2 I2 + 2 H2O → H3PO2 +4 HI (III) treatment of P4 with warm water (disproportionation) P4 + 6 H2O → PH3 + 3 H3PO2

- properties: colourless flakes, melting point: 26.5 °C, middle-strong one base acid (pKS = 1.23), acid and salts are strong reducing agents – reduce noble metals from their salts, disproportionation in warm water: 3 H3PO2 → PH3 + 2 H3PO3 (130 – 140 °C), in strong base solutions: H2PO2

- + OH- → HPO32- + H2

- application: deposition of Ni from salts on metals (pH = 4 – 6, 90 °C), plastics and other non-conductors (pH = 7 – 10/25 - 50 °C)

Salts and Organic Compounds from Phosphorus Acids (1)

Fertilizers - general importance

hypophosphoric acid

phosphinic acid

Organyl derivates are known from both tautomeric forms .

hypophosphoric acid

phosphinic acid


phosphoric acid

phosphonic acid


phosphoric acid

phosphonic acid


- Plants need for growing not only light, air, warmth and water, but additionally S, P, N, K, Ca, Mg and Fe

- Harvesting removes especially N, K, P and Ca (no back mineralisation can occur returning the minerals to the soil, only Fe is present in excess)

- Minerals can be re-added by fertilising, especially K, N and P. Solubility of the minerals in water decides about availability for the plants: high solubility → fast availability, but rapid washing out, low solubility → low, but continuous availability → “long time supply”

Manufacture of inorganic phosphates and other fertiliser salts

- phosphates/phosphites/fertilizer sulphates: neutralisation of NaOH, KOH, CaO or NH3 with phosphorus acid (phosphates), phosphorous acid (phosphites) or sulphuric acid (sulphates), precipitation and metathesis reactions

- diphosphates and polyphosphates thermal treatment of phosphate mixtures (condensation - e.g. 2 Na2HPO4 → Na2H2P2O7 + H2O), reaction time and temperature (250 - 900 °C) control polymerisation degree

Salts mainly used for fertilisation

- phosphates: superphosphates, obtained by treatment of Ca3(PO4)2 (low solubility) with 50 % H2SO4 = mixture of Ca(H2PO4)2 (high solubility) and CaSO4 superphosphate → 16 – 22 % P2O5, double superphosphate → 35 % P2O5, triple superphosphate → > 46 % P2O5 ammonium phosphates (made by neutralisation and thermal condensation), used purely and in H2O solution (high solubility) “Rhenania phosphates” – made by sintering of apatite with silica and Na2CO3 or NaOH (29 % P2O5, low solubility) “Melt phosphates” – made by melting apatite with Mg compounds and silica (21 % P2O5, low solubility) “Thomas phosphates” = slag from smelting P containing iron ores (10-18 % P2O5, low solubility)

- ammonium nitrate/ammonium sulphate (made by neutralisation) - urea (made by CO2 + NH3 → NH2COONH4 → OC(NH2)2 + H2O) - potassium chloride (mining), sulphate and nitrate (KCl + H2SO4/HNO3 or nitrates)

Salts and Organic Compounds from Phosphorus Acids (2)

Non-fertilizer applications of inorganic phosphates - sodium phosphates (general)

→ metal cleaning, phosphatising, boiler water treatment, buffer systems, food production, nutritional supplement in animal feedstuffs

- disodium dihydrogen phosphate, calcium phosphates → baking powder (additionally Ca3(PO4)2)

- tetrasodium phosphate → industrial cleaning agent

- sodium polyphosphates → added to reconstituted cheese, condensed milk, sausages, used for stabilisation of pigment suspensions and in leather tanning

- ammonium phosphates → fire protection, intumescent paints, animal feedstuffs

- tetrapotassium diphosphate → liquid cleaners

- calcium phosphates → nutritional supplement, baking powder, cleaning agent in toothpastes

NOTE: Use of phosphates for cleaning applications is decreased, because of anthropogenic phosphate entries to natural rivers and lakes causes

euthrophication.

Organic Derivates from Phosphorus Acids - phosphoric acid triesters

→ flame –retarding plasticiser, hydraulic fluids, anti-foaming agents, stabilisators

- phosphorus (V) ester acids → anti-static agent, cleaning agents, dishwasher liquids

- thiophosphoric acid derivates → herbicides (e.g. Malathion, Parathion), Zn salts are additives for lubricant oils

- amino-methylene phosphonic acid and hydroxy-ethane diphosphonic acid → detergent additives to prevent Ca precipitation in washers

- aromatic phosphorous acid esters → antioxidants, stabilisers in plastics, rubber and lubricant oils

- aliphatic phosphorous acid esters → starting materials for insecticides and veterinary products

Other Industrial Important Phosphorus Compounds (1)

Phosphorus halogen compounds - PX3, P2X4, PX5, POX3, PSX3 (X = F, Cl, Br, I), P4F6, P6Cl6, P6Br6 - mixed halogen compounds and partial substitution of halogens by hydrogen are

possible - technical importance: PCl3, PCl5 and POCl3 - synthesis:

PCl3 direct conversion of white phosphorus with dry chlorine ¼ P4 + 1.5 Cl2 → PCl3 + 320 kJ/mol PCl5 chlorine addition to PCl3, PCl3 +Cl2 → PCl5 + 124 kJ/mol POCl3 oxidation of PCl3 with oxygen at 50 – 60 °C PCl3 +1/2 O2 → POCl3 + 277.6 kJ/mol, synthesis from P4O10 and PCl3 P4O10 + 6 PCl3 + 6 Cl2 → 10 POCl3

PSCl3 PCl3 +S → PSCl3 in autoclaves at 180 °C - properties:

PCl3 colourless, smoking, toxic liquid with stabbing odour, hydrolysis in water to H3PO4 and HCl, Lewis base properties, moderate reducing agent PCl5 green-white toxic solid, decomposes at higher temperatures into PCl3 and Cl2, in the presence of water via POCl3 + HCl to H3PO4 and HCl, Lewis acid POCl3 colourless, smoking, toxic liquid PSCl3 colourless liquid, melting point: -35 °C, boiling point: 125 °C, decomposes with water to H3PO4, HCl and H2S

- use: PCl3 synthesis of H3PO3 (10 %) and alkyl substituted phosphates (detergents), of PCl5, POCl3 (33 %) and PSCl3, starting material for ligand compounds in metal-organic chemistry PCl5 chlorination agent in organic chemistry POCl3 manufacture of POX derivates (X = -OR, -NHR, -R) used as lubricant additives, softeners, flame inhibitors and insecticides PSCl3 manufacture of thiophosphoric acid ester chlorides for the production of pesticides

Other Industrial Important Phosphorus Compounds (2)

Phosphorus hydrogen compounds (phosphanes) - PH3, P2H4, P3H5 (linear compounds), P5H5, P7H3 (cyclic compounds)

with P as the electronegative partner - PH3 – exothermic compound, all other PxHy are endothermic compounds - Technical synthesises of PH3:

(1) P4 + 3 NaOH + 3 H2O → PH3 + 3 NaH2PO2 (2) 2 P4 + 12 H2O → 5 PH3 + 3 H3PO4

- Properties of PH3: colourless, toxic gas with garlic odour, low solubility in water, neutral reaction (pKB (PH3/PH4

+)= 27, pKS (PH3/PH2-) = 29),

salts: phosphides P3- (wide variety of higher phosphides from PxHy) decomposes at higher temperatures into the elements, stronger reducing agent than NH3

- Use: Manufacturing of light emitting diodes, doping of silicium, organic synthesises,

important substance in metal-organic chemistry of complexes

Phosphorus sulphur and selenium compounds - composition: P4Sn (n = 2 – 10), PSn (phosphorus polysulphides),

different thiophosphates (linear compounds), P4Sen (n = 3-5), P2Se5

- manufacture: fusing of P4 and S/Se, exothermic reactions; seperation by extraction with CS2, com- and disproportionation reactions

- application: P4S10 has some technical importance (flotation agent, lubricant additive, manufacture of insecticides)

Phosphorus nitrogen compounds

- polymeric (NPCl2)n Substitution of Cl atoms by organic groups (-OR, -NR2, -R) gives polymers with properties between caoutchouc and high severity. These polymers are used for fibres, textiles, foils and hoses. (NP-O-CH2-CF3)n polymers are used in surgery for artificial organs and chirurgical threads.

Arsenic, Antimony and Bismuth

Arsenic - natural sources: primarily sulphidic and arsenidic ores, rarely oxidic ores and in

elementary form (often mixed with antimony) - two modifications:

(I) grey rough “metallic” arsenic As 8 - stable form, conducts electricity (semi-conductor) (II) black antimony As 8 - amorphous As modification, electrical isolator, stable until 270 °C (III) yellow non-metallic arsenic As 4 electrical isolator, stable only at low temperatures and in the absence of light, converts to grey arsenic in the presence of light even at –180 °C

- As compounds are essentially in very low, but high toxic in higher concentrations (As(III) - the poison of the middle age)

- used for metal alloys especially with copper and lead (letter metals, lead accumulators), for electronic pieces (alloys with Ga and In), and in pesticides

Antimony

- natural sources: primarily sulphidic ores, rarely oxidic ores and in elementary form (often mixed with arsenic)

- two modifications: (I) grey rough “metallic” antimony Sb - stable form, conducts electricity (semi-conductor)

(II) black antimony Sb8 - amorphous Sb modification, electrical isolator, stable only below 0 °C

- Sb compounds are high toxic (similar to arsenic) - used for metal alloys especially with tin and lead to increase roughness and for

electronic pieces Bismuth

- natural sources: sulphidic and oxidic ores - non-toxic rough semi-metal (not essential for biological processes) - used for metal alloys with low melting temperatures (< 100 °C),

e.g. applied in electrical fuses


Chapter 8

Carbon Group


Group Members Carbon (C)

Silicon (Si)

Germanium (Ge)

Tin (Sn)

Lead (Pb)


12.01 28.09 72.59 118.69 207.2

Discovery unknown 1824 Berzelius 1886 Winkler

known since antique

unknown


0.02 26.3 1.4 * 10-4

2 * 10-4 1.2 * 10-5

density [g/cm³] 3.51 (diamond)

2.26 (graphite)

2.32 5.32 7.28 11.34

melting point [°C]

1410 937.4 231.9 327.4

boiling point [°C]

2250 (sublimation)

2477 2830 2687 1751


colourless diamond , polymeric black graphite, fullarenes

dark-grey, hard, brittle metal

grey-white, very brittle metal

silver-white, very soft metal

blue-grey, very soft metal

Electron negativity

2.5 1.8 1.8 1.8 1.8


-4…+4 -4…+4 -4…+4 -4…+4 +4 (+2)


oxidising power

reducing power


non-metallic (diamond),

semimetallic (graphite)

semimetallic semimetallic metallic metallic

Electrical conductivity

isolator (diamond),

semi -conductor (graphite)

semi -conductor semi -conductor conductor conductor


acid acid amphoteric amphoteric primarily

base

Stability of valence states –4 +2 +4

Physiology

base of life essentiell not essential, non-toxic

essential toxic

Electron configuration: s²p2(d10) – need of accepting or loosing

4 electrons for full saturation of electron shells

General Properties of Carbon Group Elements

Valence states: - - 4 with electropositive elements (known of all elements,

stability decreases with increasing period number) - + 2 and + 4 (+4 is favoured for C, Si, Ge and Sn, + 2 for Pb)

Hydrogen compounds and metal salts (-ides)

- formation of monomeric (XH4), polymeric (XnH2n+2 – C, Si, Ge) and cyclic (C, Si) hydrogen compounds

- stability of hydrogen compounds decreases CH4 (∆BH = – 75 kJ/mol) > SiH4 (+34 kJ/mol) > GeH4 (+91 kJ/mol) > SnH4 (+163 kJ/mol) > PbH4 (+278 kJ/mol – observed only in traces)

- bonding energy H-X decreases CH4 (416 kJ/mol) > SiH4 (323 kJ/mol) > GeH4 (289 kJ/mol) > SnH4 (253 kJ/mol)

Halogen compounds - C and Si: substitution of hydrogen atoms by halogens - Ge, Sn and Pb: XHal2 and XHal4 compounds

Oxygen compounds and binary compounds with higher chalkogenes

- C: CO and CO2 (with higher chalkogenes CY2) - Si: non-stable SiO and polymeric (SiO2)n

(with higher chalkogenes primarily SiY2) - Ge, Sn and Pb: XO and XO2 compounds

(with higher chalkogenes XY and XY2, only PbS) Acids/Bases

- C: non-stable “H2CO3” – acid properties - Si: monomeric “H4SiO4” (only stable salts) - acid properties,

condensation to polymeric acids (SiO2 · (n<2) H2O)m - Ge: “H2GeO2” and “H4GeO4” (stable only in salts or in dilution) –

acid properties - Sn: Sn(OH)2 and Sn(OH)4 – amphoteric properties

Pb: Pb(OH)2 – base reaction in H2O, plumbites in strong base solutions, Pb(OH)4 – non-soluble in water, weak amphoteric properties Pb(IV) salts in acid solutions, plumbates in base melts

Elementary Carbon

Natural sources - elementary: diamonds (Africa, Brazil, Siberia)

graphite (Madagascar, Sri Lanka, Korea, Norway etc.) - in compounds: carbonates (lithosphere, hydrosphere),

organic compounds (biosphere); carbon dioxide (atmosphere – 0.03 %, hydrosphere)

Modifications

liquid carbon

gaseouscarbon

temperature

pres

sure

graphite

diamond

diamond

graphite/diamond

graphite/melt (meta-stable)

diamond

/ melt

liquid carbon

gaseouscarbon

temperature

pres

sure

graphite

diamond

diamond

graphite/diamond

graphite/melt (meta-stable)

diamond

/ melt

diamond graphite p-T diagram of carbon

- diamond: - colourless non-metallic modification,

- electrical isolator - high hardness (used for tools) - manufacture: mining or treating of graphite at high temperatures and extreme pressures

- graphite: - grey-black semi-metallic modification with metallic brilliance - consists of hexahedron layers, connected by delocalised electrons - conducts electricity - extreme resistance against thermal stress - conversion to (synthetic) diamond at 1500-1800 °C and 53000-100000 bar (∆RH = + 1.9 kJ/mol) - manufacturing by mining of natural graphite (purification by flotation) or thermal treatment of coke, mineral oil or natural gas at 600-3000 °C

- - application: manufacturing of fire-resistant products, electrodes, paints, pencils, use as lubricant and as moderator/reflector in nuclear rectors

Special Types of Graphite

soot: low order layer structure from low temperature treatment (400-600 °C) “synthetic” electro graphite: high order layer structure from high temperature treatment (2600-3000 °C) carbon fibres: - obtained by pyrolysis of polyacrylnitrile at 2500-3000 °C, - high tensile strength and elasticity activated carbon: - microcrystalline, highly porous graphite (inner surface > 1000 m²/g) - obtained by activation of carbon with steam, air or CO2 at 700-900 °C (“burning of pores”) fullerenes: - Cn clusters, n = 60, 70, 76, 78, 84, 90, 94, …, - “footballs”, formed by side-by-side connected pentagons, - yellow-brown crystals with lower density than graphite - formed by vaporisation and rapid cooling of graphite - stable in air and water

Fullarene-60

Hydrocarbons and Halogenated Hydrocarbons

straight and branched aliphatic compounds with only single σ−σ C-C bonds (saturated hydrocarbons)

methane ethane propane butane isobutane C4H10methane ethane propane butane isobutane C4H10

straight and branched aliphatic compounds with multiple π−π C-C bonds (unsaturated hydrocarbons - olefins)

ethylene C2H4 propylene C3H6 butadiene C4H6 acetylene C2H2 propine C3H4ethylene C2H4 propylene C3H6 butadiene C4H6 acetylene C2H2 propine C3H4 cyclic compounds aromatic compounds

cyclohexane cyclopentadienecyclohexane cyclopentadiene benzene anthracenenaphthalenebenzene anthracenenaphthalene

subject of organic chemistry, more than 106 compounds Partially and fully halogenated hydrocarbons compounds

- manufactured by (1) substitution reactions with halogens Y2 (-HY) (2) addition reactions with halogens (Y2) or halogen hydrides (HY) to unsaturated hydrocarbons

- light halogenated hydrocarbons are used as solvents, blowing agents and in refrigerators and air conditioning systems (low reactivity) → use of Cl and Br substituted hydrocarbons is restricted by the Montreal protocol (ozone depleting effect in stratosphere)

- use in many organic synthesis (herbicides, intermediates to substitute functional groups)

Carbon Oxides

Carbon monoxide (CO)

- Manufacture:

- thermal treatment of coke with air at 1000 °C (→ Boudouard equilibrium), - laboratory scale: decomposition of formic acid by conc. H2SO4 HCOOH → CO + H2O

- Properties colourless, odourless, toxic and burnable gas, melting point: - 205.1°C, boiling point: -191.5 °C, low solubility in water (0.35 l/l H2O at 0 °C), triple bond between C and O

- Application - “synthesis gas” = CO/H2 mixtures for manufacture of

a large number of industrial chemicals - reducing agent (iron metallurgy - in-situ formation

from coke and air in the kiln) one of the most important industrial chemicals

Carbon dioxide (CO2)

- Manufacturing

(1) oxidation of coke in oxygen/air excess (2) by-product of lime manufacturing (2) treating carbonates with mineralic acids (laboratory scale)

- Properties colourless, non-burnable gas with acid odour,

sublimation at – 78.5°C (1 bar), liquefaction only at higher pressures (5.3 - 76.3 bar)

low solubility in water (0.9 l/l H2O at 20 °C) – acid reaction, poor reactivity, weak oxidation agent

in concentrations > 5 % toxic for humans and animals, essentially for plants

- Application inert gas, blowing agent, freezing agent (“dry ice”), neutralisation agent, sparkling agent in soft drinks

Industrial Carbon Oxide Chemistry (1)

Resources - mineral oil: - mixture of middle-heavy hydrocarbons, fractionated by

rectification (light gasoline – 30-100 °C, heavy gasoline – 100-200 °C, light oil – 200-250 °C, diesel, heavy oil – 250-350 °C, tar - >350 °C)

- Natural gas: - CH4 (80 %), C2H6, C3H8, C4H10, C5H12, impurities of H2S, CO2, N2 and He

- Coal: - complex mixture consisting of a large amount of organic (primarily poly-aromatic) compounds, contains C, O, H, N and S,

- brown coal: 65-75 % C, stone coal: 75-90 % C, anthracite: > 90 % C

Gasification of coal – Boudouard equilibrium

- Reactions

- Equilibrium plot

temperaturetemperature - conversion of coal is performed at 1000 °C in Winckler generators - product mixture = “generator gas” (70 % N2, 25 % CO and 4 % CO2)


Synthesis Gas

- Composition: - mixture of N2, CO, CO2, H2 and H2O, ratio depends on demand and is controlled by temperature, - “Water gas” = 50 % H2, 40 % CO, 5 % CO2, 4-5 % N2, traces of CH4

- Equilibrium Reactions:

temperature

- Manufacture (for details see ammonia synthesis – Chapter 7): (1) Steam reforming of natural gas, naphtha, water

2 C2H2n+2 + n H2O → n CO + 2(n+1) H2 (2) Partial Oxidation of heavy fuel oil 2 C2H2n+2 + n O2 → 2n CO + 2(n+1) H2

(3) Coal gasification C + H2O CO + H2, 2 C + O2 2 CO,


Synthesis Gas

- Application:

coal

min. oil/nat. gas

synthesis gas(syngas)

ammonia

synthesisurea resins

polyols

acetates

paraffins

olefins

alcohols

methanolsynthesis

poly-merisation

metha-nisation

polymethylene

oxo-aldehydes

synthesisoxo-

oxo-alcohols

oxidationformaldehyde

Mobile process

+ isobutene

fermentation

acetic acidanhydride (ESA)

ethylene

aromatic comp.

olefins

fuels

methyl-tert-buthyl-ether (MTBE)

process

coal

min. oil/nat. gas

synthesis gas(syngas)

ammonia

synthesisurea resins

polyols

acetates

paraffins

olefins

alcohols

methanolsynthesis

poly-merisation

metha-nisation

polymethylene

oxo-aldehydes

synthesisoxo-

oxo-alcohols

oxidationformaldehyde

Mobile process

+ isobutene

fermentation

acetic acidanhydride (ESA)

ethylene

aromatic comp.

olefins

fuels

methyl-tert-buthyl-ether (MTBE)

process

Further Important Carbon Compounds

Carbides - compounds with electropositive elements (anions C4-, C2

2-) - formed at 2000 °C

from the elements, from element compounds (especially oxides) + carbon, from element + hydrocarbon and from element compounds+ hydrocarbon

- ionic (saltlike), covalent and metallic carbides - saltlike carbides MC2: hydrolysis to acytelene (e.g. CaC2) - covalent carbides MC: high thermal resistance and hardness,

structures and properties similar to diamond (e.g. SiC, boron carbides)

- metallic carbides: with C and transition metals (IVb-VIb groups) high thermal resistance (melting points of 3000- 4000 °C), hardness similar to diamond,

conduct electricity, metallic brilliance

Hydrogen cyanide (HCN), Cyanides Manufacture

(HCN): (I) Degussa process (1200 °C, Pt catalyst) CH4 + NH3 → HCN + 3 H2

(II) Andrussow process (1200 °C, 2 bar, Pt/Rh cat.) CH4 + NH3 + 1.5 O2 → HCN + 3 H2O

cyanides: neutralisation with HCN Properties (HCN): inflammable and high toxic gas,

soluble in water (weak acid reaction, Ks = 2.1 · 10-9), complexing agent

Application: galvanisation (salts), gold mining and extraction, processes in organic chemistry,

e.g. methyl methacrylate (acid and salts) Phosgene (COCl2)

Manufacture: CO + Cl2 → COCl2 at 720-750 °C Properties: highly reactive and high toxic gas

Application: synthesis of poly-urethanes, chlorination agent for metal oxides (e.g. SnO2 → SnCl4)

Carbon disulfide CS2

Manufacture: (I) C + S2 → CS2 at 720-750 °C (II) CH4 + 2 S2 → CS2 + 2 H2S at 650-750 °C Properties: endothermic liquid, inflammable and high toxic

Application: viscose industry (rayon), cellophane production, synthesis of CCl4, production of vulcanisation accelerators, flotation agents, corrosion inhibitors, herbicides and pharmaceuticals

Elementary Silicon (1)

Natural sources - second most common elements - occurs only in compounds - minerals: quartz sand (SiO2), silicates

Manufacturing of the metal

- metallurgical grade Si: thermal treating of quartz with coke in an electrical furnace (at 2000 °C for 1-2 h, ∆RH = 690.4 kJ/mol) SiO2 + C → SiO + CO SiO + 2 C → SiC + CO 2 SiC + SiO2 → 2 Si + 2 CO

- electronic grade Si: - purification of metallurgical grade Si by (1) conversion into SiHCl3 (300 °C) Si + 3 HCl → SiHCl3 + H2 (2) Distillation of SiHCl3 (3) decomposition of SiHCl3 at 1000 °C yielding to highly purified Si SiHCl3 + H2 → Si + 3 HCl - formation of singly crystals by zone melting - zone melting in presences of traces of volatile compounds of the dopant or by thermo-neutron bombardment (Si → P)

-

Elementary Silicon (2)

Properties

- dark-grey, hard, brittle metal, semi-conductor - lattice structure similar to diamond - metallurgical grade Si: 98.5-99.7 % - electronic grade Si: >99.999 % - reacts with electronegative elements only at high temperatures (exothermal

reactions, but passivation, with O2 at 1000 °C, with N2 at 1400 °C, with S at 600 °C, with C at 2000 °C, only with F2 at room temperature)

Application

- component of steel - special alloys with iron (ferrosilicon, 8-13 % Si, 87-95 % Fe) - alloys with Al, Cu and Ti - semiconductor components (diodes, transistors, electronic circles, processors, solar

cells)

Hydrogen Silicon Compounds

straight and branched aliphatic compounds with only single σ−σ Si-Si bonds (saturated silanes)

monosilane disilane trisilane tetrasilane iso-tetrasilane polysilane H3Si-(SiH2 )n-SiH3

nmonosilane disilane trisilane tetrasilane iso-tetrasilane polysilane H3Si-(SiH2 )n-SiH3

n

cyclic compounds (Si > 4)

cyclopentasilane cyclohexasilanecyclopentasilane cyclohexasilane → no Si multiple bonds (no formation of π-π bonds) → no stable unsaturated or aromatic compounds

wide variety of compounds, similar reactions to organic chemistry

Monosilane SiH4 :

- manufacture: (1) decomposition of Mg2Si with acids in the absence of air Mg2Si + 4 H+ → SiH4 + 2 Mg2+ (2) hydrogenation of SiCl4 with LiH in molten LiCl/KCl at 400 °C SiCl4 + 4 H- → SiH4 + 4 Cl-

- properties: - endothermic colourless gas - melting point: -184.7 °C, boiling point: -112.3 °C - stable up to 300 °C in the absence of air, than decomposes to Si and H2

- under air: inflammable, SiH4 + 2 O2 → SiO2 + 2 H2O + 1518 kJ/mol - in the presence of water: SiH4 + 2 H2O → SiO2 + 4 H2 + 374 kJ/mol - with halogens and hydrogen halogenides stepwise replacing of H by Hal

- application: manufacture of ultra pure Si

Technical Important Binary Silicon Compounds

Silicon halides: - SiF4 manufacture: CaF2 + H2SO4 → 2 HF + CaSO4

SiO2 + 4 HF → SiF4↑ + 2 H2O (in conc. H2SO4) properties: - highly exothermic compound, stable under dry air, decomposes in water: 3 SiF4 + 2 H2O → SiO2 (aq) + 2 H2SiF6 application: hydrolysis → HF manufacture

- SiCl4 : manufacture: Si + 4 HCl → SiCl4 + 2 H2/ Si + 3 HCl → SiHCl3 + H2 (300 °C)

properties: colourless, smoking liquid with sticking odour application: - manufacture of electronic grade Si

- synthesis of organic silicon compounds - siliconisation of metallic surfaces

- manufacture of highly dispersed SiO 2

Silicon dioxide (SiO 2) – quartz - manufacture: - mining of quartz sands, re-crystallisation in H2O - - flame hydrolysis of SiF 4 and SiCl4 - properties: - high stable, exothermic solid

- crystal and glass (amorphous) modification - reacts only with HF at room temperature,

- reacts with alkali hydroxides in molten state - application: - glass industry, foundries, chemical apparatus

- manufacturing of silicates, of enamels, of ceramics and of SiC

- polishing agent, inorganic filler - “piezoelectrical” effect – use in quartz watches

Silicon nitride (Si3N4)

- manufacture: (1) from the elements at 1100-1400 °C 3 Si + 2 N2 → Si3N4 + 750 kJ/mol (2) 3 SiO2 + 2 N2 + 6 C → Si3N4 + 6 CO (3) 3 SiCl4 + 4 NH3 → Si3N4 + 12 HCl

- properties: - colourless solid with high hardness but low density

- stable up to 1900 °C - resistant against corrosion and mechanical stress

- application: - special ceramics - manufacturing of chemical apparatus, high quality mechanical tools - mechanical and motor engineering - fittings

Silicon carbide (SiC): see carbides under “Further Important Carbon Compounds”

Metal Silicides: metallic hard materials

Silicon Acid and Silicates

Manufacture: - Silicates: melting of alkali carbonates or hydroxides

with quartz sand - Acid: (1) solving alkali silicates in H2O and

precipitation of the acid by slow acidification (2) hydrolysis of monomeric SiX4 compounds

SiX4 + 4 H2O → “H4SiO4” + 4 HX Properties: - weak acid, meta-stable only in dilution

- condensation = tri-dimensional polymerisation (1) initial reaction

Polymerisation goes via micro-particles (sol) to a highly viscous

silicon acid/water mixture (gel) to crystalline SiO2.

(2) secondary structure

chains bands layers

- determining structure by blocking of polymerisation sides by metal

cations - base unit: SiO4 tetrahedrons, coupled by corners, edges and surfaces

Application: - glass and alkali silicates - zeolithes (with AlO4

- salts) - glass fibres

- cement (CaSiO3) - natural and synthetic fillers

Glassware

Starting materials: - quartz sand - soda ash and potash, Glaubers’s salt

- lime - lead oxide

- borax - kaolin and feldspar (Al)

Manufacture: melting of the mixture of starting materials at 1200-1650 °C and stepwise careful cooling with homogenisation

Properties: - amorphous mixed silicates, “freezed melt”

- high thermal and chemical resistance - good electrical isolator - softening temperature

between 550-650 °C (soda-lime glass) and 2000 °C (quartz)

Composition of typical glassware:

Zeolithes

Manufacture: - from natural zeolithes (e.g. kaolin) - conversion by shock-heating at > 550 °C,

followed by suspension in NaOH solution at 70-100 °C, product: zeolithe A (ion exchanger in for detergents)

- synthetically by common precipitation from Na water glass and NaAl(OH)4 solutions

at high temperatures and partially high pressures - synthetically by sol-gel techniques

(in alcohols with metal alcoxides as initiators)

Structure: - consisting of (SiO4) and [AlO4]- tetrahedrons with large, but specific cavities

- cations are delocalised in the cavities and can be exchanged by other cations

Elementary cell of sodalithe Tertiary structure of zeolithe A

Application: - ion exchangers (e.g. in detergents) - molsieve (after thermal treatment at 400-550 °C

to remove water from the cavities) - specific adsorption agent

- catalyst (cracking and isomeristaion of hydrocarbons) and catalyst support (high specific surface areas)

Silicones (1)

Manufacturing:

Precursors: - (chloro)methylsilanes - (chloro)phenylsilanes

- obtained by Rochow Müller process - (300/500 °C, CuO catalyst,

ZnO as an activator) Si + 2 CH3Cl → (CH3)2SiCl2

Si + 2 PhCl → Ph2-SiCl2

Polymerisation: - hydrolysis of products in 25 % HCl at 100 °C

gives cyclic and linear siloxanes (1:1 – 1:2) - ring opening with KOH or with strong minaralic

acids (H2SO4 )

Silicon oils

Molecular structure: linear polysiloxanes

Properties: - thermal stability (300 °C) - viscosity only poor dependent from

temperature - high electrical resistance - low surface tension - odourless, tasteless, physiologically inert

Application: - heat transfer media, lubricants, hydraulic oils, transformer oils,

brake fluids, paint flow improvers, gloss improvers, defoaming agents, mold releasing agents, component of skin creams and protective polishes

Silicones (2)

Silicone Rubbers

- crosslinked polysiloxanes

- application: sealing compounds in the construction industry, in sanitary sector, glass sector and automobile industry, adhesives for heat-resistant bonds and seals

Silicon Resins

- poly-organosiloxanes with a high portion of branched tri- or tetrafunctional siloxy groups

- thermal stable, weather resistant, hydrophobic - application: electrically insulating lacquers,

corrosion protection lacquers (pigmented with zinc dust), stoving enamels, coil coating of metallic plates for facades, rendering plastics scratch resistant

Germanium

Natural sources - rather seldom sulphidic minerals (not in technical use)


- outgoing from waste gases of Zn manufacture - Process steps:

(1) extraction GeO2 and ZnO of fly ash with H2SO4 (2) precipitation of oxides at pH = 5 with NaOH (3) conversion of oxides to chlorides with HCl (4) distillation → separation of GeCl4 and ZnCl2 (5) Hydrolysis of GeCl4 to GeO2 (6) reduction to Ge with H2

- high purity Ge (electrical grade) is obtained by zone melting

Properties of the metal - grey-white, very brittle metal with semi-conductor properties - stable under air, in water, in base solutions and in non-oxidising acids - transformation to GeO2 by concentrated H2SO4 and HNO3

Application of the metal

- transistors - optical lenses, prisms, windows (high IR transparence) - special alloys and superconductors

Germanium compounds

- typical reactions and compounds of IVa group elements - oxidation state + 4 is favoured compared to + 2 (both are stable) - amphoteric (predominantly acid) character of hydroxides - no technical importance of single compounds

Tin

Natural sources - occurs primarily in form of sulphidic and oxidc ores - minor amounts of elementary metal


- thermal reduction with coke 360 kJ/mol + SnO2 + C → Sn + 2 CO

- recycling of tinplate wastes in an electrochemical process

Properties of the metal - silver-white, very soft metal with conductor properties - essential, non-toxic - high stability under air (reacts only in the heat) and in water - oxidised by hot strong base solutions (forming stannates(IV))

and by concentrated acids (forming Sn(II) salts

Application of the metal - dishes - corrosion inhibitor for iron sheet metals by impregnation with molten Sn

(forming tinplate) - solder tin (alloy with 30-60 % lead, eutectic point)

Tin compounds

- typical reactions and compounds of IVa group elements - oxidation state + 4 is favoured compared to + 2 (both are stable) - amphoteric (predominantly acid) character of hydroxides

- SnCl4: - obtained from the elements by treating of tinplate wastes

with chlorine: Sn + 2 Cl2 → SnCl4 + 511.6 kJ/mol - colourless smoking liquid, Lewis acid properties - used as homogeneous Friedels-Crafts catalyst and for synthesis of organic Sn compounds

- SnO2: - white pigment for glazes and enamels - organic Sn compounds:

- partially high toxicity - use as PVC stabilisator, for vulcanisation of silicones, as biocides and anti-fouling paints

Lead

Natural sources - occurs primarily in form of sulphidic ores


- “roast reduction process” PbS + 1.5 O2 → Pb + SO2 PbO + CO → Pb + CO2

- “roast reaction process” 3 PbS + 3 O2 → PbS + PbO + 2 SO2 PbS + 2PbO → 3 Pb + SO2

- purification by melting with air or by electrochemical process (enrichment of silver impurities)

- recycling of accumulators

Properties of the metal - blue-grey, very soft metal with conductor properties - non-essential, but high toxic - high stability under air (passivation, reacts only in the heat)

oxidised in oxygen-containing water - oxidised by base solutions (forming plumbites(II))

and by acids (forming Pb(II) salts, but passivation by low-soluble salts – PbSO4, PbCl2, PbF2)

- absorption of radioactivity -

Application of the metal - tanks for strong corrosive

chemicals - accumulators - liquid in heating baths - protection against radioactivity - alloys with antimony – high

mechanical strength – used for bearings

charging

dis-chargingenergy

electrons

charging

dis-chargingenergy

electrons

Lead compounds

- typical reactions and compounds of IVa group elements - oxidation state + 2 is favoured compared to + 4

(strong oxidation agent) - amphoteric (predominantly acid) character of hydroxides

- PbO · PbO2: - used in glass manufacture, corrosion inhibiting paint - lead salts: - oxides and chromate are used in oil paints


Chapter 9

Earth Metals (Boron Group)


Group Members Boron (B)

Aluminium (Al)

Gallium (Ga)

Indium (In)

Thallium (Th)


10.81 26.98 69.72 114.82 204.37

Discovery 1808 Gay-Lussac,

Thenard, Davy

1825 Oerstedt

1875 de Boisbaudran

1863 Reich, Richter

1861 Crookes


0.001 7.7 1.6 * 10-3

1 * 10-5 5 * 10-5

density [g/cm³] 2.46

2.70 5.91 7.31 11.85

melting point [°C]

660.3 29.8 156.6 303.5

boiling point [°C]

2250 (sublimation)

2330 2403 2070 1453

Electron negativity

2.0 1.5 1.6 1.7 1.8


-1,+1,+3, complex anions

+1/+3 +1/+3 +1/+3 +1 (+3)


non-metallic metallic metallic metallic metallic

Electrical conductivity

isolator conductor conductor conductor conductor


acid amphoteric amphoteric amphoteric base

Stability of valence states +1 +3

Physiology

not essential, non-toxic

not essential, non-toxic

- - toxic

Electron configuration: s²p1(d10) – need of loosing 3 electrons

for full saturation of electron shells

General Properties of Earth Metals

Valence states:

- + 1 and + 3 (+3 is favoured for B, Al, Ga and In, + 1 for Tl) - complex anions of boron in compounds with electropositive elements


- covalent compounds

- formation of dimeric B2H6 and of polymeric (AlH3)n - GaH3, InH3, TlH3 were not found

- stability of hydrogen compounds: ½ B2H6 (∆BH = +18 kJ/mol) < 1/n (AlH3)n (-11 kJ/mol)

Halogen compounds

- B: BHal3, B2Hal4 and (BHal)n compounds - Al: polymeric compounds - Ga, In, Tl: XHal, X-XHal4 and XHal3 compounds

Binary compounds with chalkogenes

- X2Y3, Ga, In, Tl forms additionally mono-chalkogenides X2Y

Binary compounds with pnictogenes

- XY – hard compounds, diamond-like structure, partially semiconductors

Acids/Bases

- B: H3BO3 – acid properties - Al, Ga, In: amphoteric X(OH)3 compounds - Tl: weak base Tl(OH)3, strong base TlOH

Elementary Boron

Natural sources - occurs only in compounds - minerals: kernite Na2B4O7 · 4 H2O and borax Na2B4O7 · 10 H2O

Manufacture of the element

- amorphous boron (technical grade) thermal reduction of B2O3 with Mg B2O3 + 3 Mg → 2 B + 3 MgO + 533 kJ, purification by treating with boiling diluted HCl

- crystallised boron of high purity (1) reduction of boron halogenides (Cl, Br) with hydrogen (1000 - 1400 °C, W or Ta catalyst) 2 BHal3 + 3 H2 → 2 B + 6 HHal

(2) thermal decomposition of BI3 at 800-1000 °C 2 BI3 → 2 B + 3 I2 + 142 kJ

(3) thermal decomposition of B2H6 (600 - 800 °C, BN, W or Ta catalyst) B2H6 → 2 B+ 3 H2 + 36 kJ

Properties of the element - properties between metals and non-metals - one glass like amorphous and four crystalline modifications - complex crystal structures with B12 icosa-hedrons as base unit,

partially including hetero atoms (e.g., B24C, B24N, B12P)

B12 icosa-hedron

- stable up to 400 °C, reacts at >700 °C with air, at >400 °C with Cl2, at >700 °C with S and at 900 °C with N2

- stable in non-oxidising acids, in oxidising acids up to 250 °C - melting with alkali yields to alkali borates - favoured oxidation state: +3

Boron Compounds (1)

Hydrogen compounds - base element BH3 is not stable (electron shortage) and

has strong Lewis acid properties - stabilisation

by intra-molecular adduct formation (homologous rows BnHn+2, BnHn+4, BnHn+6, BnHn+8, BnHn+10), as anions (e.g. BH4

-, B3H82-),

by formation of adducts - toxic compounds with sickening odour, inflammable, short boranes are not stable in

water - stability and acid strength increase with number of B atoms

(e.g. B6H10 - weak acid < B4H10 < B10H14 < B18H22 - strong acid), - technical importance: B2H6 (made from 2 BF3 + NaH → B2H6 + NaF)

for hydroboration reactions in organic chemistry NaBH4 (made by Schlesinger process at 250-270 °C – B(OMe)3 + 4 NaH → NaBH4 + 3 NaOMe or from borax, Na and H2 – Na2B4O7 · 7 SiO2 + 16 Na + 8 H2 → 4 NaBH4 + 7 Na2SiO3

Halogen compounds

- BHal3, B2Hal4 and (BHal)n compounds - BF3: - manufactured by treating borates with fluorspar and conc. H2SO4

B2O3 + 6 HF → 2 BF3 + 3 H2O - colourless, sticking gas, strong Lewis acid - application: Friedels-Crafts catalyst, flowing agent

- HBF4:- manufactured by treating boron acid with hydrogen fluoride H3BO3 + 4 HF → HBF4 + 3 H2O - use: strong acid for reactions catalysed by protons, galvanisation

- BCl3: - manufacture: B2O3 + 3 C + 3 Cl2 → 2 BCl3 + 3 CO - colourless, smoking gas, high reactivity with water (to H3BO3 + HCl) - use: semiconductor industry, Friedels-Crafts catalyst, high purity boron manufacture

Oxygen compounds and acids

- B2O3, HBO2 and H3BO3 - borax - Na2B4O7 · 10 H2O – used for ceramics, enamels, glassware - perborates - NaBO2 · 2 H2O – used in detergents as leaching agent

Boron Compounds (2)

Boron nitride BN - manufacture: B2O3 + 2 NH3 → 2 BN + 3 H2O (800-1200 °C) or

B2O3 + 3 C + N2 → 2 BN + 3 CO - properties: highly inert material - use: high temperature lubricant,

lining of rocket burning chambers, plasma burners and nuclear reactors

Boron carbide B4C

- manufacture: B2O3 + 7 C → B4C + 6 CO (at 2400 °C) or 2 B2O3 + C + 6 Mg(4 Al) → B4C + 6 MgO + 2 Al 2O3

- properties: highly inert and hard material (similar than diamond) - use: abrasive, manufacture of metal borides, armour plates,

neutron catcher in nuclear reactors

Elementary Aluminium

Natural sources - occurs only in compounds - minerals: corundum Al 2O3, hydrargillite (Al(OH)3,

feldspar, clays, bauxite (alumosilicates), cryolithe Na3[AlF6]

Manufacture of the metal - Bayer process for Al 2O3 manufacture

(1) treating of bauxite with 35-38 % NaOH at 140-250 °C and 5-7 bar for 6-8 hours Al(OH)3 + NaOH → Na[Al(OH)4] – separation from Fe(OH)3 (2) precipitation by dilution (decrease of pH)

(3) calcinations to form Al 2O3

- electrolysis of cryolyth (82 %)-alumina (18 %) mixture at 940-980 °C

liquid Al

molten electrolyte

carbon cathode

carbon anode

isolation

liquid Al

molten electrolyte

carbon cathode

carbon anode

isolation

Properties of the metal - light silver-white and elastic metal - high affinity to oxygen, but passivation - not stable in acids (forming salts and H2) and

bases (forming aluminates and H2) - favoured oxidation state: +3

Application of the metal

- vehicles and aircraft - containers and packaging - construction industry - office and household equipment - iron and steel industry (alloys) - aluminothermal welding (3 Fe3O4 + 8 Al → 4 Al2O3 + 9 Fe + 3341 kJ) - important compounds: Al 2O3, Al(OH)3, Al2(SO4)3, AlCl3, NaAlO2, AlF3, NaAlF6,

spinells


Chapter 10

Alkaline Earth Metals


Group Members

Beryllium Be

Magnesium Mg

Calcium Ca

Strontium Sr

Barium Ba

Radium Ra

Atom Number 4 12 20 38 56 88 Rel. Atomic Mass

9.01 24.31 40.08 87.62 137.33 [226] radioactive

Discovery 1828 Wöhler

1809 Davy

1808 Berzelius,

Pontin

1790 Grawford

1774 Scheele

1898 Curie


2.7 * 10-4 2.0 3.4

3.6 * 10-2 4 * 10-2 1 * 10-10

melting point [°C]

1278 648.8 839 768 710 ~ 700

boiling point [°C]

~ 2500 1105 1482 1380 1537 ~ 1140

Density [g/cm³] at 25 °C and 1 bar

1.85 1.74 1.54 2.63 3.65 unknown

Electron negativity

1.5 1.2 1.0 1.0 0.9 0.9


+2 +2 +2 +2 +2 +2


moderate reducing

agent

strong reducing

agent


silver-white or yellow-white metals


amphoteric

base

Physiology highly toxic essential essential not essential, not toxic

not essential, but toxic

Electron configuration: s² (d10) – need of loosing 2 electrons

for full saturation of electron shells General properties: - occur in oxidation state +2

- non-noble metals, forming mostly exothermic compounds with primary ionic character (NOTE: Be forms covalent compounds!)

Chemical Properties of Alkaline Earth Metals


- monomeric XH2 compounds (except Be)

- stability of hydrogen compounds decreases

BeH2 (∆BH ~ 0 kJ/mol, cannot be obtained from the elements < MgH2 (∆BH = -74 kJ/mol) < CaH2 (∆BH = -184 kJ/mol) > SrH2 (∆BH = -177 kJ/mol) > baH2 (∆BH = -172 kJ/mol)

- salt-like hydrides (except “BeH2”), stable under air

- reaction with water to hydroxides and H2

Halogen compounds - Be: covalent halogen compounds (BeHal2)n

with Lewis acid character

- higher elements: ionic salts XHal2 Chalkogen compounds

- Be: BeY, stable in air and water

- higher elements: XY, stable in air, hydrolysis in water (forming hydroxides and H2Y)

Acides/Bases - Be: Be(OH)2, soluble in acid and base solutions - higher elements: X(OH)2, only soluble in acids, base reaction

Beryllium (1)

Natural sources - rather rare element

- minerals: beryl - 3 BeO · Al2O3 · 6 SiO2, bertrandite - 4 BeO · 2 SiO2 · H2O

- deposits in USA, Russia, Argentina and Brazil Manufacture of the metal

Pre-processing: - extraction of minerals with H2SO4 - separation of aluminium salts with by precipitation with (NH4)2SO4 - precipitation of Be(OH)2 with NH3

(1) - treating of Be(OH)2 with NH4HF2 forming (NH4)2BeF4 - thermal decomposition of (NH4)2BeF4 at 900 – 1000 °C (NH4)2BeF4 → BeF2 + 2 NH3 + 2 HF (recycling of NH3 and HF) - chemical reduction of BeF2 with Mg: BeF2 + Mg → Be + MgF2 at 1300 °C

(2) - thermal conversion of Be(OH)2 to BeO - BeO + C + Cl2 → BeCl2 + CO at 800 °C - purification of BeCl2 by distillation at 485 °C - electrolysis of molten, water-free BeCl2: BeCl2 → Be + Cl2 at 350 °C

Properties

- grey, hard, brittle metal, stable under air up to 600 °C and in water - low density - solved by diluted acids forming H2 (Be + 2 H+ → Be2+ + H2),

passivation by oxidising acids - reacts with electronegative elements only in the heat - high toxicity of the pure metal (dust) and of Be compounds

Application

- limited by high price and high toxicity - manufacturing of Be-Cu alloys for electrical equipment

(0.5 – 2.5 % Be – increase of mechanical strength) - moderator and reflector material in nuclear plants - aerospace applications

Beryllium (2)

Chemistry - formation of primarily covalent compounds - “electron shortage compounds” with Lewis acid character ,

→ stabilisation by adduct formation and complexes

→ polymerisation by Lewis acid-base interactions

- amphoteric character of Be(OH)2

Be(H2O)42+ Be(OH)2 Be(OH)4

2- soluble in

acid solutions precipitation

in neutral solutions soluble in

base solutions

- similarity to aluminium covalent, polymeric hydrogen compounds (BeH2)x and (AlH3)y, Lewis acid properties of halogenides, amphoteric character of hydroxides

Magnesium (1)

Natural sources - eighth most frequent element

- minerals (examples): magnesite – MgCO3, dolomite – CaCO3 · MgCO3, carnallite – KCl · MgCl2 · 6 H2O, kieserite – MgSO4 · H2O, asbestos (silicates), olivine – [Mg, Fe)2SiO4]

- deposits in China, Russia, North Korea, Brazil, Australia

- remarkable amounts in seawater Manufacture of the metal

(1) Pre-processing: thermal conversion of MgO MgO + C + Cl2 → MgCl2 + CO + 153 kJ/mol,

electrolysis of molten, water-free mixture of MgCl2 (8 - 24 %) and other alkali and alkaline earth metal chlorides, MgCl2 → Mg + Cl2 at 700 - 800 °C, removal of molten Mg (favoured process, 80 % of world Mg)

(2) thermal reduction of dolomite with ferrosilicon (vacuum, 1200 °C) 2 (CaO · MgO) + Si(Fe) → 2 Mg + Ca2SiO4 (slag) + Fe (slag) Properties

- silver, middle-hard metal, oxidised under air and in water, but passivated - low density - solved by diluted acids forming H2 (Mg + 2 H+ → Mg2+ + H2),

passivation by oxidising acids - strong reducing agent - reacts with electronegative elements strongly exothermically (bright light) after

activation in the heat - formation of compounds with intermediate ionic covalent character - essential element

Application

- lightest construction metal - manufacturing of Al-Mg alloys (< 10% Mg, casting alloys, wrought alloys) and Mg

based alloys with Al, Mn, Zn, Si, Be (Mg > 90 %, motor industry) - reducing agent in organic and inorganic chemistry, Grignard reactions - desulphurisation and deoxidification agent in iron and steel industry - pyrotechnical applications - manufacturing of metals (Be, Ti – Kroll process)

Magnesium (2)

Important magnesium salts Magnesium carbonate (MgCO3) - magnesite

- Manufacturing: mining of natural sources, purification by gravitational separation, flotation or magnetic separation; synthetic salt produced by precipitation from Mg salts Mg2+ + (NH4)2CO3 → MgCO3↓ + 2 NH4

+ or by carbonating of MgO (e.g. obtained from dolomite) MgO + 2 CO2 + H2O → Mg(HCO3)2 in aq. solution

- Application: manufacturing of MgO, thermal insulating material, filler in paper, plastics and rubber, additive in table salt and in pharmaceuticals

Basic magnesium carbonate (Mg(OH)2 · 4 MgCO3 · 4 H2O)

- Manufacturing: calcination of Mg(HCO3)2 - Application: mild neutralisation agent,

used in medicine to neutralise antacid

Magnesium hydroxide (Mg(OH)2) - Manufacturing: precipitation from aqueous solutions of Mg salts

Mg2+ + Ca(OH)2 → Mg(OH)2 + Ca2+ - Properties: low solubility in water, basic character,

soluble in acids and in NH4+ containing solutions

Magnesium oxide (MgO)

- Manufacturing: (1) calcination of magnesite or dolomite at > 550 °C MgCO3 → MgO + CO2 (2) precipitation from brines and seawater with limestone Mg2+ + Ca(OH)2 → Mg(OH)2 + Ca2+, calcination of Mg(OH)2

Mg(OH)2 → MgO + H2O temperature: 600-900 °C - “caustic” MgO, 1600-2000 °C - “sintered” MgO, melting at 2800-3000 °C – “fused”/”dead burnt” MgO - Application: caustic MgO – fertilizers, animal feedstuff,

building materials, chemical and pharmaceutical products, water treatment sintered MgO – refractory industry (isolation of metallurgic kilns) fused MgO – insulating material in electrical heating

Magnesium (3)

Important magnesium salts (continuation) Magnesium Chloride (MgCl2)

- Manufacturing: (1) from brines and seawater Dow Chemical process - precipitation of Mg(OH)2 with lime - Mg(OH)2 + Ca(OH)2 + 2 HCl + 2 H2SO4 → MgCl2 + CaSO4↓ + 4 H2O - evaporation at 200 °C → MgCl2 · 2 H2O - evaporation at 300 °C → MgCl2 (2) from Mg carbonate or oxide (burnt magnesite) Norsk-Hydro process (magnesite + seawater) 2 MgO + 2 Cl2 + (2) C → 2 MgCl2 + CO2 (or CO) (3) MPLC process

MgO + Cl2 + CO → 2 MgCl2 + CO2 (4) dehydration of hexahydrate

- Application: electrochemical manufacture of Mg (80 %), granulation of fertilizers, as a dust binder, in oil and sugar industry, mixtures of MgO and MgCl2 used for production of Sorel cement and lightweight building panels

Magnesium sulphate (MgSO4)

- Manufacturing: (1) mining of kieserite or from brines (2) byproduct of K salt processing (3) reacting of carbonates or seawater with H2SO4

- Application: fertilizer (80 %), manufacture of K2SO4, Na2SO4 and K-Mg sulphate (potash magnesia), textile and cellulose industries, production of building materials, refractory materials, animal feedstuffs and motor oil additives

Calcium, Strontium, Barium and Radium

Common chemical properties

- metals are oxidised in the presence of air and water - solved by diluted acids forming H2 - strong exothermic reactions with electronegative elements after thermal activation - formation of primarily ionic compounds with cations M2+ - low solubility of sulphates, fluorides, carbonates, silicates and phosphates (but high

solubility of hydrogen carbonates and monohydrogen/dihydrogen phosphates)

- low solubility of hydroxides, basic reaction of the solution

Calcium

Natural sources - 5th most frequent element, widely distributed all over the world

- minerals: limestone – CaCO3, dolomite – CaCO3 · MgCO3

gypsum – CaSO4 · 2 H2O, anhydrite CaSO4, apatite – Ca5(PO4)3F,

Manufacture of the metal

- thermal reduction with aluminium (vacuum, 1200 °C) 6 CaO + 2 Al → 3 CaO · Al2O3 (cement slag) + 3 Ca (g)

- electrochemical processes are no longer operated

Properties of the metal

- silver-white metal with low density Application of the metal

- reducing agent in manufacturing of Zr, Th, U and rare earth metals - refining agent in metallurgy - maintenance-free batteries (Pb/Ca alloys) - manufacturing of SmCo5 magnetic materials

Limestone and Construction Materials

Calcium carbonate (CaCO3) - limestone - manufacturing: open cast mining,

synthetic fine-grained CaCO3 by carbonating milk of lime (pigments for paint and paper industry)

Calcium sulphate (gypsum – CaSO4 · 2 H2O, anhydrite CaSO4)

- manufacturing: open cast mining, by-product of manufacture of H3PO4 and of HF, by-product from waste gas desulphurisation

The “limestone cycle” in manufacture of construction materials

Ca(OH)2(slake lime/lime hydrate)

CaO(quicklime)

CaCO3(limestone)

CaCO3(limestone)

(1) (3)(2) Ca(OH)2(slake lime/lime hydrate)

CaO(quicklime)

CaCO3(limestone)

CaCO3(limestone)

(1) (3)(2)

(1) calcinations of limestone at 1000 – 1200 °C 178.4 kJ/mol + CaCO3 → CaO + CO2 (2) slaking of CaO = slow addition of water CaO + H2O → Ca(OH)2 + 65.2 kJ/mol (3) use of slake lime in construction, binding of slake lime with atmospheric carbon dioxide Ca(OH)2 + CO2 → CaCO3 + H2O cement = 3 CaO · SiO2, obtained by thermal treating of a mixture of CaO, SiO 2 and some additives (Fe2O3, Al2O3) at 1450 °C lime mortar = mixture of slake lime and sand gypsum mortar = suspension of anhydrite, hardening by local solution and re-crystallisation Other applications of limestone, quicklime and slake lime

limestone: - fertilizer - steel industry

quicklime and slake lime: - metallurgy (removal of P and S) - water and effluent treatment - chemicals (carbides, cyanamides, Na2CO3 – Solvay process) - agriculture and sugar industry - refractory materials - flue gas desulphurisation

Other Calcium Compounds

Calcium carbide (CaC2)

- manufacturing: reacting highly purified CaO with coke in an electrical furnace at 2000 – 2200 °C 464 kJ/mol + CaO + 3 C → CaC2 + CO

- application: - formation of acetylene CaC2 + 2 H2O → C2H2 + Ca(OH)2 (exothermic !), widely applied in welding and in manufacture of special cast iron, - formation of calcium cyan amide (CaCN2) - desuphurisation and deoxidation of raw iron and steel

Calcium cyan amide (CaCN2) - nitro-lime

- manufacturing: CaC2 + N2 → CaCN2 + C + 296 kJ/mol (700–900 °C) - application: - long-time NH3 fertilizer

(CaCN2 + 3 H2O → CaCO3 + 2 NH3 + 91.3 kJ/mol) - herbicide

- manufacturing of organic chemicals Calcium fluoride (CaF2)

- manufacturing: mining and purification of fluorspar - application: - manufacturing of HF

- enamel industry - optical prisms and lenses (high UV transmission)

Calcium fluoride (CaCl2)

- manufacturing: waste product from many processes

- soda manufacturing (Solvay process) CaCO3 + 2 NaCl → Na2CO3 + CaCl2 - propylene oxide from chlorhydrin - treating waste HCl with limestone vaccum and atmospheric pressure evaporation, traded as 30 – 45 % solution or as 75 % flakes

- application: - dust binder (road re-construction, mines) - cooling, defrosting and antifreeze agent (e.g. road de-icing in strong winters – main use) - drying agent (laboratory scale) → available amount exceeds demand

Strontium and its Compounds

Natural sources - minerals: celestine – SrSO4,

strontianite – SrCO3

- deposits in Mexico, Spain, Turkey and Great Britain Manufacture of the metal

- elementary metal is not used in technical scale - laboratory scale: chemical and electrochemical reaction

Properties of the metal

- light gold-yellow metal with low density - vaporised Sr emits red light

Strontium compounds

SrCO3: - manufacturing: mining - application: manufacture of special glasses (CRT-screen glassware for colour TV and computer monitors), magnetic materials, pigments and fillers, electrolytic Zn manufacture (precipitation of Pb and Cd salts)

Sr(NO3)2: - manufacturing: SrCO3 + 2 HNO3 → Sr(NO3)2 + CO2 + H2O - application: fireworks

Barium and its Compounds

Natural sources - minerals: heavy spar/ barite – BaSO4,

(witerite – BaCO3, not mined) - deposits in China, USA, India, Russia

Manufacture of the metal

- elementary metal is only used in special applications (getter material in the manufacture of valves)

- laboratory scale: chemical and electrochemical reaction Properties of the metal

- gold-yellow metal with low density - vaporised Ba emits green light

Barium compounds

BaSO4: - manufacturing: (1) mining (2) oxidation of BaS with Na2SO4 - application: drilling-mud for oil and gas exploration (90 % of mined BaSO4), white pigment (manufacture of paper, paint, rubber and plastics)

BaCO3: - manufacturing: 3 step process (1) BaSO4 + 4 C → 4 BaS + 4 CO (rotary kiln, 1200 °C) (2) BaS + CO2 + H2O → BaCO3↓ + H2S or BaS + Na2CO3 → BaCO3↓ + Na2S (in aq. solution) - application: tile and ceramic industry (preventing bleading of Na and Ca sulphates), special ceramics as Ba ferrite and Ba titanate, glass industry – producing of special optical glassware and CRT-screens, manufacturing of photographic papers

BaO2: - manufacturing: glowing of BaCO3 and coke BaCO3 + C → BaO + 2 CO, thermal oxidation of BaO at 500-600 °C 2 BaO + O2 → 2 BaO2

- application: igniting agent Ba(NO3)2: manufacturing: BaCO3 + 2 HNO3 → Ba(NO3)2 + CO2 + H2O

- application: fireworks


Chapter 11

Alkali Metals


Group Members

Lithium (Li)

Sodium (Na)

Potassium (K)

Rubidium (Rb)

Caesium (Cs)

Francium (Fr)

Atom Number 3 11 19 37 55 87 Rel. Atomic Mass

6.94 22.99 39.10 85.47 132.91 [223] (radio-active)

Discovery 1817/18, Arfevedson,

Davy

1803, Davy

1807, Davy

1861/62, Bunsen,

Kirchhoff

1860, Bunsen,

Kirchhoff

1939, Perey


2.0 · 10-3 2.7 2.4 9.0 · 10-3 3.0 · 10-4 1.3 · 10-21

melting point [°C]

180.5 97.8 63.6 38.9 28.4 ~ 27

boiling point [°C]

1347 881.3 753.8 688.0 678 ~ 660

Density [g/cm³] at 25 °C and 1 bar

0.53 0.97 0.86 1.5 1.9 unknown

state at 25 °C and 1 bar malleable silver metals

malleable golden metal

Electron negativity

1.0 0.9 0.8 0.8 0.7 0.7


+1 +1 +1 +1 +1 +1


strong reducing agents


typical metals


basic oxides and hydroxides

Electron configuration: s1(d10) – need of loosing 1 electron

for full saturation of electron shells

General Properties of Alkali Metals (1)

Manufacture of metals - electroylsis of molten, water-free salts (Downs process):

2 LiCl → 2 Li + Cl2 at 610 °C 2 NaCl → 2 Na + Cl2 at 600 °C

- chemical reduction: KCl + Na → K + NaCl (850 °C) 2 RbOH + Mg or Ca → Mg(OH)2 or Ca(OH)2 + Rb Cs2Cr2O7 + 2 Zr → 2 Cs + 2 ZrO2 + Cr2O3 (500 °C , high vacuum)

Physical properties

- malleable metals with low melting and boiling temperatures - low density (Li, Na and K less than water) - coloured vapours at higher temperatures

(consisting of atoms and molecules M2) (Li: red, Na: yellow, K: violet, Rb: red, Cs: blue)

Chemical propertes

- strong reducing agents - metals will be oxidised under air atmosphere even at room temperature (traces of

humidity are necessarily) - occur in nature only as salts - oxidation number in compounds: +1, - similarity between Li and Mg (e.g. solubility of salts) - formation of binary compounds with all non-metallic (electronegative) elements

(e.g. H – hydrides, C – carbides, N – nitrides, S – suphides) - Na and K are essentially for biological processes


Hydrogen compounds

- stable exothermic ionic hydrids MH with salt-like properties, stable up to 360 °C (RbH) - 970 °C (LiH)

- formation from the elements M + 1/2 H2 → MH at temperatures: LiH 600-700 °C, KH, RbH and CsH – 350 °C, NaH – 250-300 °C)

- formation enthalpies: LiH –93.2 kJ/mol, NaH –57 KJ/mol, KH –56 kJ/mol, RbH –55 kJ/mol, CsH –50 kJ/mol)

- strong reducing agents

- reactions: 2 MH + O2 → M2O + H2O (hydrides are flammable)

MH + H2O → MOH + H2 MH + NH3 → MNH2 (amides) + H2 MH + Hal2 → MHal + HHal - Application

LiH: hydrogenation agent in organic and inorganic chemistry NaH: base compound in organic reactions (Claisen condensation, aldol additions, alkylation and acylation reactions, reducing agent in inorganic chemistry, synthesis of hydride compounds, manufacturing of pure metals (Hydrimet process)

Hydroxides

- Synthesis: (1) hydroxides are made by electrolysis of chlorides

2 M+ +2 Cl- + 2 H2O → 2 MOH + 1/2 H2 + 1/2 Cl2 (NaOH/KOH: mercury and membrane process, diaphragm process only for NaOH) see chlorine alkali electrolysis, chapter 6) (2) “caustification” of carbonates – “old” industrial process Na2CO3 + Ca(OH)2 → 2 NaOH + CaCO3 – only for NaOH

- Properties: strong basic character, no decomposition to oxides in the heat

- Application: NaOH - widely used “basic substance” , e.g. for manufacture of soaps, dye stuffs and cellulose,

- for synthesis of a lot of chemicals, for purification of fats, oils and petroleum KOH - manufacture of soaps, potash, phosphates and other potassium compounds


Chalkogen compounds (general)

- formation of compounds M2Y (chalkogenides) with partially covalent character,

M2Yn>1 (perchalkogenides) and M>2O (suboxids) - properties of sulphides → chapter 7

Oxides, suboxides, peroxides and ozonides

Oxides M2O

Peroxides M2O2 Hyperoxides MO2

Ozonides MO3

Li Li2O2 Li2O2

4 Na Na2O

1 Na2O22 NaO2

3 NaO35

K K2O6 K2O2

1 KO22 KO3

5 Rb Rb2O

6 Rb2O21 RbO2

2 RbO35

Cs Cs2O6 Cs2O2

1 CsO22 CsO3

5

Manufacture: 1 – from the elements under careful oxygen and temperature control 2 – combustion of M in under oxygen excess 3 – formation under large oxygen excess and high pressure 4 – reaction of MOH with H2O2 and thermal decomposition 5 – oxidation of hyperoxides or of hydroxides with ozone at temperatures < 0 °C (MO2 + O3 → MO3 + O2, 3 MOH + 2 O3 → 2 MO3 + MOH · H2O + 1/2 O2), 6 – com-proportionation reaction M2O2 + M → 2 M2O or 2 MNO3 + 10 M → 6 M2O + N2) Properties:

- oxides are stable up to 500 °C, peroxides are stable up to 500 - 600 °C (Li2O2 up to 200 °C) hyperoxides are stable (except RbO2 only up to 450 °C) decomposition of ozonides at room temperature

- reaction with water peroxides: decomposition to H2O2, O2 and OH- hyperoxides: 2 O2

- + 2 H2O → O2 + H2O2 + 2 OH- ozonides: 4 O3

- + 2 H2O → 5 O2 + 4 OH- Application: Na2O2 – leaching agent Li2O2, Na2O2, KO2 – oxygen source and CO2 absorber in respirators (e.g. 4 KO2 + 2 CO2 → K2CO3 + 3 O2)


General properties of salts - oxidation number of alkali metals +1 - mostly ionic character (depending on the nature of the anion) - high thermal stability (depending on the nature of the anion), high melting and

boiling points - high solubility in water (mostly kg/l, except some Li salts and perchlorates of K,

Rb, Cs) dissociation into solvatisated cations and anions (solutions conduct electricity)

- NH3 acts partially as a “pseudo alkaline metal” Halogen compounds

- stable ionic halogenides MHal with salt-like properties - formation from the elements in partially strong exothermic reactions - primary natural sources of alkali metals and halogens

Lithium and Its Salts

Elementary lithium

- manufacturing: electrolysis of molten LiCl/KCl mixture at 400-460 °C - application: - synthesis of Li hydride and Li amide

- synthesis of organic lithium compounds (reducing agent, polymerisation catalysts) - manufacture of extremely light and strong Al-Li alloys (2 - 3 % Li) for space applications - batteries

Lithium carbonate (Li2CO3)

- manufacturing: precipitation from soluble Li salts 2 Li+ + CO3

2- → Li2CO3 - application: - agent for decreasing melting temperature

in aluminium manufacture - flux in glass, enamel and ceramic industries - medicine (psychiatry) - manufacture of fire-resistant glassware

Lithium hydroxide (LiOH)

- manufacturing: Li2CO3 + Ca(OH)2 → CaCO3 + 2 LiOH - application: - manufacturing of Li soaps and greases

Lithium hydride (LiH)

- manufacturing: 2 Li + H2 → 2 LiH at 700 °C - application: - drying agent

- hydrogen storage - reducing agent in organic chemistry (LiAlH4, LiBH3)

Lithium nitrate (LiNO3)

- manufacturing: Li2CO3 + 2 HNO3 → 2 LiNO3 + CO2 + H2O in aqueous solution

- application: red fireworks

Manufacture of Soda Ash - The Solvay Process Process steps

(1) preparation of a concentrated NaCl solution (2) saturation of the solution with NH3 under cooling (3) saturation of the solution with CO2 at 50 °C

NH3 + CO2 + H2O → NH4+ + HCO3

- NH4

+ + HCO3- + Na+ + Cl- → NaHCO3↓ + NH4

+ + Cl-

(4) thermal decomposition of NaHCO3 at 170 – 180 °C

2 NaHCO3 → Na2CO3 + H2O + CO2 (recycling to step 3)

(5) producing of additional CO2 by calcination of limestone at 900 °C CaCO3 → CaO + CO2

(6) regeneration of ammonia

2 NH4+ + 2 Cl - + CaO → 2 NH3 + Ca2+ 2 Cl- + H2O

(CaCl2 cannot used in further processes and is an waste difficult to depose)

Summary: 2 NaCl + CaCO3 → Na2CO3 + CaCl2

(occurs in aqueous solution in the opposite direction)

Sodium and its Salts

Salt

Manufacture Application

elementary Na

- electrolysis of molten NaCl (modified Downs process)

- reducing agent, catalyst in organic chemistry

- manufacture of NaH, NaBH4, Na2O2 etc.

- coolant in nuclear reactors (fast breeders)- sodium-sulphur

batteries NaCl - mining or underground solving of

rock salt deposits (in Germany - Staßfurt/Zielitz, Austria, Spain, USA, Russia) and purification

by flotation, - vaporising, freezing or electrodialysis

of sea water

- starting material for all other inorganic Na compounds

(e.g. Na2CO3, NaOH, Na2SO4, Na2B4O7, Na2SiO3)

- raw material for chlorine alkali electrolysis

- food industry Na2CO3

(soda ash) - mining of trona deposits (USA),

purification by solving, evaporating and calcination

- Solvay process (Europe) 2 NaCl + CaCO3 → Na2CO3 + CaCl2

- glass industry - synthesis of inorganic Na salts

- pulp and paper industry - soap and detergent production

NaHCO3 - Na2CO3 +H2O + CO2 → 2 NaHCO3 (high purity)

- food industry (baking powder production)

- animal feedstuff - rubber, chemical, pharmaceutical, textile, leather and paper industries

NaNO3 - mining of natural deposits (Chile) - Na2CO3 +2 HNO3 → 2 NaNO3 + H2O

+ CO2

- fertilizer

Na2SO4 (Glauber’s

salt)

- mining of natural deposits (Russia, USA, Canada)

- 2 NaCl + H2SO4 → Na2SO4 + HCl (800 °C)

- 2 NaCl + MgSO4 → Na2SO4 + MgCl2 (in aq. solutions)

(deep temperature precipitation of Na2SO4)

- pulp and paper industry - additive in detergents

- glass industry - chemical industry

NaHSO4 - 2 NaSO4 + H2SO4 → 2 NaHSO4 - byproduct of CrO3 manufacture

(Na2Cr2O7 + 2 H2SO4 → 2 CrO3 + 2 NaHSO4 + H2O)

- cleaning agents - flux

NaB4O7 (borax)

- extraction from borate minerals (dissolving in H2O, followed by selective crystallisation at 60 °C)

- dehydratisation by calcination at 350-400 °C

- glass, enamel, china and ceramic industries

- manufacture of perborates for detergents

- flux, falme and corrosion inhibitor

Potassium and its Salts

Salt

Manufacture Application

elementary K - KCl + Na → K + NaCl (760-880 °C, favoured process)

- 2 KF + CaC2 → CaF2 + 2 C + 2 K (1000-1100 °C)

- only limited importance - manufacture of KO2

- manufacture of low melting Na-K alloys (reducing and drying agent, heat

transfer medium) KCl - mining or underground solving of salt

deposits (deposits in Germany - Staßfurt/Zielitz, Hanover, Werra/Fulda region - France, Canada, USA, Russia)

and purification by flotation

- starting material for all other inorganic K compounds

- production of K containing fertilizers (KCl, K2SO4, KNO3)

- metallurgy, enamel industry, manufacture of soaps

- manufacture of special IR transparent optical glassware

KBr - halogenation of potash with Fe(II, III)-bromide

4 K2CO3 + Fe3Br8 → 8 KBr + Fe3O4 + 4 CO2

- bromation of potash 3 K2CO3 + 3 Br2 → 5 KBr + KBrO3 + 3

CO2

- photographic industry - manufacture of special IR transparent

optical glassware

KI - halogenation of potash with Fe(II, III)-iodide

4 K2CO3 + Fe3I8 → 8 KI + Fe3O4 + 4 CO2

- reduction of KIO3

- photographic industry - manufacture of special IR transparent

optical glassware

K2CO3 (potash)

- carbonation of KOH KOH + CO2 → K2CO3 + H2O

(precipitation, calcination at 250-350 °C)

- glass and enamel industry, pigment manufacture

- manufacturing of soaps and detergents - food industry

- starting material for many inorganic and organic K compounds

KNO3 (saltpetre)

- KCl + NaNO3 → NaCl + KNO3 - 2 KCl + 2 HNO3 + 1/2 O2 → 2 KNO3

+ Cl2 + H2O

- fertilizer - component of gun powder

K2SO4 - 2 KCl + H2SO4 → K2SO4 + HCl (700 °C)

- 2 step process in aq. solutions: (1) 2 KCl + 2 MgSO4 → K2SO4 ·

MgSO4↓ + MgCl2 (2) K2SO4 · MgSO4 + 2 KCl → 2 K2SO4

+ 2 MgCl2

- fertilizer (trading of K2SO4 and K2SO4 · MgSO4)

course on inorganic chemistry by frank klose chapter 1

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