Course on Inorganic Chemistry
by Frank Klose
Chapter 1
Elements and Compounds,
Atoms and Molecules –
Structures and Bonds
Substances, Compounds and Elements
Substances
Heterogeneous
substances
Homogeneous
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
-
1999
-
Proposal of the “Periodic Table of Elements” by
Mendelejew and Meyer (independently)
Discovery of the elements 114
(Joint Institute for Nuclear Research Dubna,
Russia), 116, 118 (Berkeley, California, USA)
Number of elements discovered
200
175
150
115
82
100
50
12 14 15
32
0
Antiquity
19th century
13th century
20th century
17th century
Theoretical maximum
18th century
Percentage of Elements
Earth’s crust
Element
Percentage
[mg/kg]
O
Si
Al
Fe
Ca
Na
K
Mg
Cl
Ti
H
P
Mn
F
Ba
Sr
S
C
N
Zr
V
Cr
467600
278600
81600
50200
36400
28400
26000
21000
18100
4400
1400
1000
950
625
425
375
260
200
20
165
135
100
Rb, Ni, Zn, Ce, Cu, Y,
La, Nd, Co, Sc, Li, Nb,
Ga, Pb, Th, B
10 - 100
Pr, Br, Sm, Gd, Ar, Yb,
Cs, Dy, Hf, Er, Be, Xe,
Ta, Sn, U, As, W, Mo,
Ge, Ho, Eu
1 - 10
Tb, I, Tl, Tm, Lu, Sb,
Cd, Bi, In
0.1 – 1
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
< 0.1
Human “biomass”
Percentage
Element
[mg/kg]
O*
C*
H*
N*
Ca*
P*
S*
K*
Na*
Cl*
Mg*
Fe *
Zn*
Si*
Rb
F*
Sr
Zr
Cu*
Br
Sn*
Nb
I*
Al
Pb
Cd
Ba
Mn*
V*
B
Se *
Mo*
As*
Co*
Cr*
Li
Ni*
*) essentially
611000
236000
94000
28000
14000
9330
2330
2270
1400
1400
440
56
40
18.7
18.7
10.7
4
4
2.67
1.87
1.87
1.33
0.933
0.467
0.467
0.4
0.267
0.267
0.267
0.187
0.187
0.0667
0.0467
0.0373
0.0267
0.0267
0.0133
The Atom and its Components
1808
1897
1913
1911
-
1921
1926
-
1932
-
Hypothesis of atoms by Dalton
Discovery of the electron by J. J. Thomson
Discovery of the proton by E. Rutherford
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
Discovery of the neutron by W. D. Harkins
Schrödinger equation, begin of the quantum
mechanical description of atoms
Atom model by W. Heisenberg using electron
orbitals
Atom
Core
Shell
Protons
Neutrons
Electrons
positive
charged
no
charge
negative
charged
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:
NA Avogadro number:
h
Planck constant:
1.6022 * 10-19 C
6.0221 * 1023
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 functio n
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
dxy orbital
d x 2 − y 2 orbital
px orbital
dxz orbital
py orbital
pz orbital
dvz orbital
d z 2 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
7d
Energy
7p
n=7
7s
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 (s 2, p6, d 10, f14) have the highest stability. Half
occupied d subshells (d 5) 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
Periodicity - The Size of Atom Orbitals
electron
core
2nd shell
1st shell
electron
core
3rd shell
2nd shell
1st shell
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)
Term scheme for sodium (Na)
→ using of outer electron transitions
(∆l = ±1, ∆j = 0, ±1,
s → p and p → d transitions)
→ specific for each element
→ laser technology
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
(1)
primary
X-ray radiatation
(2)
(3)
secondary
X-ray
radiatation
electron energy [eV]
absorption of primary X-ray
radiatation
→ remove of a inner
electron → ionisation
transfer of an electron from
an outer shell to the leak
emission of secondary X-ray
radiation (specific for the
element)
The Periodic Table of Elements (PTE)
Main group elements
n
Ia
IIa
IIIa
d block elements
(transition metals)
IIIb
IVb
Vb
VIb
VIIb
VIIIb
Ib
IVa
Va
VIa
VIIa
VIIIa
IIb
Lanthanides
Actinides
s block
(l = 0)
d block
(l = 2)
p block
(l = 1)
Ia – VIIIa, Ib – VIIIb = group numbers
1-7 = numbers of periods
atomic number Z
(= number of protons)
rel. atomic mass
(= molecular mass)
electron
negativity
chemical symbol
name of the element
colour = metallic or non-metallic character or
acid-base properties of the oxides
Prediction and Discovery of Germanium
18691871
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
relative atom mass
colour
density [g/cm³]
specific heat capacity [J/gK]
melting point [°C]
valency
oxide
formula
density [g/cm³]
acid/base
properties
chloride
formula
density [g/cm³]
boiling point [°C]
ethyl compound formula
density [g/cm³]
boiling point [°C]
Properties
Proposed by
Mendelejew
1871
Properties found
by Winkler
1886
Present State
72
dark grey
5.5
0.306
high
4
AO2
4.7
predominantly
acid
ACl4
1.9
60 - 100
A(C 2 H5 )4
0.96
160
72.32
grey white
5.47
0.318
4
GeO 2
4.703
acknowledged
72.61
grey white
5.32
0.310
937.4
4
GeO 2
4.228
acknowledged
GeCl4
1.887
86
Ge(C 2 H5 )4
0.99
163
GeCl4
1.8443
83.1
Ge(C 2 H5 )4
0.991
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 Quantum mechanical LCAO-MO-theory
Lennard-Jones
1927-31
Heitler,
Quantum mechanical “valence bond theory”
London, Slater,
Pauling
1931
Pauling
Hybridisation
Types of Chemical Bonds (1)
Octet rule:
The electron configuration of noble gases (s 2, s 2p6, s 2p6d10,
s 2p6d10f14 – 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)s 2, ((n-1)d5(n)s 2 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°
Types of Chemical Bonds (2)
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
polarity of the ions
polarisation of the covalent bond
There exists a continuum between covalent and ionic parts of bonds!
Molecule
LiF
LiCl
LiBr
LiI
CsCl
BaO
Ionic part of the
bond
0.87
0.73
0.59
0.55
0.75
0.43
Molecule
NO
CO
HCl
HBr
HI
H2
Ionic part of the
bond
0.015
0.01
0.18
0.12
0.05
0
Types of Chemical Bonds (3)
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, P 4, 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. SO42-: 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 + O 2 → 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
- number of bonds =
number of bonded MO - number of anti-bonded MO
δ orbital
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
sp
sp²
Involved
atom orbitals
Geometric
form
Type of the
molecule
s, px
s, px , py
linear
triangle
sp³
s, px , py , pz
tetrahedral
AB2
AB3
AB2
AB4
AB3
sp²d
sp³d
s, px , py , dxy
s, px , py , pz,
quadratic
triangle
bipyramidal
(2 tetrahydrons)
quadratic
pyramid
dz2
sp³d
s, px , py , pz,
dx 2 −y 2
sp³d²
s, px , py , pz,
dz2 , dx 2 −y 2
octahedron
AB2
AB4
AB5
AB3
AB2
AB5
AB6
AB5
AB4
Geometry of
the molecule
Example
linear
triangle
V form
tetrahedral
triangle
pyramid
V form
quadratic
triangle
bipyramidal
T form
linear
quadratic
pyramid
BeCl2
BF3
SO2
CH4
NH3
octahedron
quadratic
pyramid
quadratic
SF6
BrF5
H2 O
XeF4
PF5
ClF 3
XeF2
BrF5
XeF4
Special Cases of Hybrid Orbitals
Ethylene
(C-C double bonds)
Acetylene
(C-C triple bonds)
= σ bond + π bond
(sp2 hybrid orbitals)
= σ bond + 2 π bonds
(sp hybrid orbitals)
Diborane B2H6
(“electron shortage compounds”)
Benzene
(“aromatic systems”)
2 electrons triple center bond
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. SiO 2).
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, Jür gen 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
ability to be oxidised
electron affinity/negativity, ionisation energy
metallic character
non-metallic character
basic strenght of oxides
acid strenght of oxides
Ia
IIa
IIIa
IIIb
electron affinity,
ionisation energy,
non-metallic character,
acid strenght of oxides,
oxidation state (valency)
Lanthanides
Actinides
IVb Vb
VIb
VIIb
VIIIb
Ib
IIb
IVa Va
VIa VIIa VIIIa
Course on Inorganic Chemistry
Chapter 2
Chemical Reactions
The Chemical Equilibrium
Consider the reaction
k
α A + β B →
γ C + δ D with r1 = k1 * [A]α * [B]β
1
If back reaction
γ
δ
k
γ C + δ D 
→ α A + β B with r-1 = k-1 * [C] * [D]
−1
also occurs, we have a chemical equilibrium described by
K=
k1
[C ] γ ⋅ [ D] δ
=
k −1 [ A ] α ⋅ [ B ] β
“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)
energy
activated complex
k1 = k0, 1 * exp (-EA, 1/RT)
catalyst
k-1 = k0, -1 * exp (-EA, -1 /RT)
EA, 1 ≠ EA, -1
In case of equilibrium
r1 = r-1 ≠ 0
A+B
C+D
→ dynamic equilibrium
reaction coordinate
Special cases:
(1)
Nernst’s distribution law
(2)
Henry Dalton’s law
(3)
K=
K′ =
cA , phase1
cA , phase2
cA , liqiuid solution
K
=
RT
p A , gas phase
[ B] + ⋅ [ A] −
electrolytic dissociation
Kc =
[ AB]
-4
(K c < 10 - 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
→
Decrease of pressure
→
(3)
favours the reaction with ∆rV < 0
favours the reaction with ∆rV > 0
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
Oxidation
Reducing
agent
Reduction
Oxidising
agent
+ electrons
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
HCl
H2O
±0
Oxidation number of hydrogen +1
Oxidation number of chlorine – 1
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
Zn pole
Cu pole
CuSO4 solution
ZnSO4 solution
membrane
Galvanic cell (voluntary)
Anode (negative pole - oxidation):
Zn → Zn2+ + 2 eCathode (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: ε = ε0 +
R ⋅T
c
⋅ lg Ox .
Z ⋅F
c Re d .
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 solution
b) non-metallic elements and compounds
acid 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
K = 10-14, pH = -log [H3O+]
OH- + H3O+
Acid anhydrides = compounds (oxides or metal cations) forming Brönstedt acids
first by the reaction with water
e.g.
SO3 + 2 H2O
Al3+ + 7 H2O
H2SO4 + H2O
HSO4- + H3O+
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+ , Fe 2+
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:
soft acids:
hard bases:
soft bases:
cations with small diameters, high positive charge and no non-bonded
electrons,
→ H+ , cations from s 1, s 2, s 2p1 and d10s2p2 elements
→ forming mainly ionic bonds
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
anions with a central atom highly charged
and possessing a high electronegativity
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).
Literature/References for Figures
(1)
Arnold Frederik Holleman, Egon Wiberg,
Lehrbuch der anorganischen Chemie
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
Course on Inorganic Chemistry
Chapter 3
Noble Gases
Overview About the Group
Group
Members
Atom
Number
Rel. Atomic
Mass
Helium
(He)
2
Neon
(Ne)
10
Argon
(Ar)
18
Krypton
(Kr)
36
Xenon
(Xe)
54
Radon (Rn)
4.00
20.18
39.95
83.80
131.29
Discovery
1895
Ramsey
1898
Ramsey
1894
Ramsey,
Rayleigh
1898
Ramsey
1898
Ramsey
[222]
(radioactive)
1900
0.000524
0.00182
0.9340
0.000114
0.000087
Percentage
in air
[Vol.-%]
86
Dorn,
Rutherford,
Soddy
6 * 10-18
Electron configuration:
s²p 6(d 10) – 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 24α species)
Physical and Chemical Properties
Group
Members
melting point
[°C]
boiling point
[°C]
vaporization
enthalpy
[kJ/mol]
1st ionisation
energy [eV]
Helium
(He)
-272.1
(2.5 MPa)
-268.9
(4 He)
0.092
Neon
(Ne)
-248.6
Argon
(Ar)
-189.4
Krypton
(Kr)
-156.6
Xenon
(Xe)
-111.5
Radon
(Rn)
-71
-246.0
-185.9
-152.9
-107.1
-61.8
1.86
6.28
9.68
13.70
18.02
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”)
- XePtF 6 by Barlett (1962, theoretically predicted by Pauling 1933)
Known noble gas compounds
- RnF 2, fluorides, oxides and oxifluorides of Xe, chlorides of Xe, KrF 2
- no compounds of He, Ne , Ar
The Air Rectification Process by Linde
Joule Thomson effect:
Gases can be cooled by adiabatic expansion, if
temperature δ a is less than inversion temperature and
µJT (Joule Thomson coefficient) is positive.
Joule Thomson parameter
and inversion temperature for
different gases
The Linde process
compressed air
(heating)
cooler
expanded air
cross flow
heat exchanger
air inlet
(cooling)
δ A, pA
throttle valve
(as the „ideal gas“)
compressor
liqiud air
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:
δ A - δ E = µJT * (pA – pE) *(273.15/(273.15 + δ A))²
Fractions of the technical rectification
→ further purification in additional rectification steps
δE, pE
Halogen Compounds of Noble Gases
Xe + F2
C , Ni tube
400
°
→
XeF2
(colourless solid)
Xe + 2 F2
C , 0.6 MPa, Xe / F2 = 1: 5
400
°
   
→
XeF4
(colourless solid)
Xe + 3F2
250 °C , 5 MPa, Xe / F2 = 1: 20
200
−
    
→
XeF6
(colourless solid)
Kr + F2
183 ° C , 20 mbar
−
 
→
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:
-
all noble gas halogen compounds have strong oxidation power
§ XeF2 : Cl- → Cl2 , IO 3 - → IO 4 , BrO 3 - → BrO 4 - (all in aqueous solutions),
Fluorination of NO2 to FNO2 , reaction with F2 to XeF4 and XeF6
§ XeF4: Pt → PtF 4 , Hg → Hg2 F2
§ XeF6 : Hg → HgF 2 , AuF 3 → Au(V)
§ KrF2 : ClF 3 → ClF 5 , Ag → AgF 2 , Hg → HgF 2 ,
[KrF]+ strongest known oxidation agent
7 KrF2 + 2 Au → 2 [KrF][AuF 6 ] → AuF 5 + Kr + F2
- increases with increasing atomic number of noble gas
atom (RnF 2 /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))
Oxygen Containing Compounds of Noble Gases
→ only known compounds:
XeO 3 , XeO 4 , H4 XeO 6 ,
XeOF2 , XeO 2 F2 , XeOF4 , XeO3 F2 , XeO2 F4
Molecular structures of XeO 3 and XeO 4
Xenon(VI)-oxide (XeO 3 )
- preparation:
XeF6 + 3 H2 O → XeO 3 + 6 HF
3 XeF4 + 6 H2 O → Xe + XeO 3 + 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 (XeO 4 )
- preparation:
basic hydrolysis of XeO 3
XeO 3 + OH- → HXeO 4 2 HXeO 4 - + 2 OH- → XeO 46-+ Xe + O2 +2 H2 O
- properties:
XeO 4 – yellow liquid (< - 40 °C)/colourless gas,
XeO 4 6- yellow solutions
XeO 4 – explosive above – 40 °C
strong oxidation power (ClO 3 - → ClO 4-, Cr3+ → Cr2O7-, Mn2+ →
MnO4 -,(IO 3 - → IO 4-)
Oxiflouride Compounds
- preparation:
- properties:
deep temperature hydrolysis of XeF4 ,
reaction of xenon fluorides with xenon oxides
colourless crystals, which can be hydrolysed,
poor stability
Literature/References for Figures
(1)
Arnold Frederik Holleman, Egon Wiberg,
Lehrbuch der anorga nischen Chemie
101st edition, Berlin [u.a.] : de Gruyter, 1995
(4)
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
Course on Inorganic Chemistry
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 H2 O)
Hydrogen isotopes
Name
atom core
composition
rel. atomic
mass
natural
percentage
protium H
1 proton
1.0078
deuterium D
1 proton +
1 neutron
2.0141
tritium T
1 proton +
2 neutrons
3.0160
99.9855 %
0.0145 %
10-15 %
s1
→ needs to spent or to accept one electron
→ occurs in elementary form as diatomic H2
Electron configuration:
ortho and para hydrogen
spins of protons
electron shell
atom cores
para-hydrogen
para-hydrogen
ortho -hydrogen
percentage of ortho-hydrogen [%]
percentage of para-hydrogen [%]
ortho-hydrogen
- 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, c p,
vapour pressures)
absolute temperature
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 → CaC 2 + 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 pressure
solution
phase
hydride
phase
mixed phase
(solution + hydride)
plateau region
hydrogen uptake (mol H/mol metal)
-
by reaction of halogenides with LiH, NaBH2 or LiAlH4
EHal n + 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
Hydrogen metal
ratio x of the
compound EHx
Stability increase
IIIb
=31
IVb
=2
Vb
=2
VIb
=22
VIIb
n.d.
VIIIb
=22
Hydrogenation
catalysts
1 – including lanthanoides and actinoides
2 – only known from Cr, Ni (at high pressures) and Pd
n.d. – not discovered
Ib
1
IIb
2
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)
-
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 2 nd 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
period
-
-
-
period
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-1well 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 > Cl 2 >Br 2 …, 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 + HAlH3 + H-
BH4AlH4-
Heavy and Super-heavy Water
rel. molecular mass
density (25 °C) [g/cm³]
maximum density [g/cm³] /
Temperature of density
maximum [°C]
melting point [°C]
boiling point [°C]
dissociation constant pKW (25
°C)
Toxicity
H2 O
“light water”
D2 O
“heavy water”
18.02
0.997
1.000/3.98
20.03
1.104
1.106/11.23
T2 O
“super-heavy
water”
22.03
1.214
1.215/13.4
0.000
100.00
14.000
3.81
101.42
14.869
4.48
101.51
15.215
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)
Literature/References for Figures
(1)
Arnold Frederik Holleman, Egon Wiberg,
Lehrbuch der anorganischen Chemie
101st edition, Berlin [u.a.] : de Gruyter, 1995
(5)
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
Course on Inorganic Chemistry
Chapter 5
Halogens
Overview About the Group
Group
Members
Atom Number
Rel. Atomic
Mass
Discovery
Fluorine
(F)
9
19.00
Chlorine
(Cl)
17
35.45
Bromine
(Br)
35
79.9
Iodine
(I)
53
126.90
Astatine
(At)
85
209.99
1886
Moissan
1774
Scheele
1826
Balard
1811
Courtois
0.11
6 * 10-4
5 * 10-5
-101.00
-7.25
113.60
1940
Corson,
McKenzie,
Segré
3 * 10-24
(radioactive)
300
-34.06
58.78
185.24
335
yellowgreen
gas
dark brown
liquid
solid
4.0
3.0
2.8
violet
crystals
with
metallic
brilliance
2.5
-1
-1…+7
-1…+7
-1…+7
-1…+7
Percentage on
0.06
earth [Mass%]
melting point
-219.62
[°C]
boiling point
-188.13
[°C]
state at room colourless weak
temperature
yellow gas
(25 °C)
and 1 bar
Electron
negativity
valence
numbers in
compounds
Reducing
Power
Oxidation
power
Electron configuration:
2.2
s²p 5(d 10) – 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 – Na3 AlF6
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 (H2 O + 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: KMgCl 3 · 6 H2O,
MgCl 2 · 6 H2O,
KMgCl(SO4) · 3 H2O
Manufacturing in industry:
1. Electrolysis of NaCl brines
(2 H2O + 2 NaCl → H2 + 2 NaOH + Cl 2)
Restrictions to the process:
- prevention of formation of hypochlorite in the solution
(2 OH- + Cl 2 → OCl - + Cl -)
- suppressing of contact between H2 and Cl 2
(→ 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 + Cl 2
3. Thermal oxidation of hydrogen chloride (Deacon process)
4 HCl + O2
2 H2O + 2 HCl + 2 Cl 2
catalyst:
CuCl 2 (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 Cl 2 + 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 state
The Diaphragm Process for Manufacturing Chlorine
-
-
General:
separation of chloride oxidation and hydrogen reduction
by an asbestos membrane
Reactions:
Anode:
Cl→
0.5 Cl 2 + e - (ε 0= 1.36 V)
Na + + OHNaOH
Cathode:
H2O + e
→
H2 + OH(ε 0 = 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 Cl 2 (but permeability for dissolved Cl 2)
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 Cl 2 + e - (ε 0= 1.36 V)
Na + + OHNaOH
Cathode:
H2O + e
→
H2 + OH(ε 0 = 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 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
Manufacturing in 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- + BrO 3- + 3 H2O (a)
5 Br- + BrO 3- + 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
(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
of
water
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
Manufacturing in industry:
1.
from residual solutions of Chilean niter production (50 %)
acidification of brines with H2 SO3
(treatment with gaseous SO2) – reduction of IO3- to I(HIO3 + 3 H2 SO3 → 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 Cl 2/H2 SO4 → “blow out” of iodine with air
→ purification by reduction with SO2 and re-oxidation with Cl 2 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, PF6solubility 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
Hydrogen
fluoride
Hydrogen
chloride
Hydrogen
bromide
Hydrogen
iodide
-271
-92
-36
+26
+3.05/+2.87
+1.63/+0.42
+1.06/+1.06
+0.54/+0.54
state at room
temperature
(25 °C)
and 1 bar
colourless gas
with
sticking smell,
toxic
colourless gas
with
sticking smell,
toxic
colourless gas
with
sticking smell
toxic
melting point
[°C]
boiling point
[°C]
-83
-114
-87
colourless
furning liquid
with
sticking smell
toxic
-35
+20
-85
-67
+26
hydrofluoric
acid
hydrochloric
acid
hydrobromic
acid
hydroiodic acid
+3.2
unlimited
-6.1
507
-8.9
612
-9.3
425
Name of the
poor
compound
Formation
enthalpy ∆B H0
[kJ/mol]
ε 0 (2 X-/X2)
[V] (pH=0/14)
Reducing
power
Oxidation
power
Name of the
aqueous
solution
pK s
solubility
[l/l H2O]
Manufacturing and Use of Hydrogen Fluoride
Manufacturing in industry:
Conversion of fluorspar CaF 2 (Bayer Process)
CaF 2 + 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 MF x and acid salts MF x · (HF)n
Application:
- manufacture of inorganic fluorides (AlF 3, BF 3, UF 6, 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 (SiO 2 + 4 HF → SiF 4 (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
→ Al2 O3 + 3 H2 O (300-400 °C)
Al2 O3 + 6 HF → 2 AlF 3 + 3 H2 O (400-600 °C)
Chemie Linz AG process
2 Al(OH)3 + H2 SiF 6 → 2 AlF 3 + SiO 2 + 4 H2 O (100 °C)
- use:
flux in the aluminium industry
Sodium Aluminum Hexafluoride (Cryolite Na3 AlF 6 )
- manufacture: 6 NH4 F + 3 NaOH + 2 Al(OH)3
→ Na3 AlF 6 + 6 NH3 + 6 H2 O
- use:
electrolytic manufacture of aluminium
Alkali Fluorides (NaF, KHF 2 , NH4 F · HF)
- manufacture: NaOH + HF or H2 SiF 6
- use:
NaF – water fluoridation
KHF 2 – frosting agent in glass industry,
synthesis of F2
NH4 F – oil extraction
Hexafluorosilicates (M2 SiF 6 )
- manufacture: 2 MCl + H2 SiF 6 → M2 SiF 6 + 2 HF
- use:
wood protection
Na2 SiF 6 - water fluoridation
Uranum Hexafluoride (UF 6 )
- manufacture: UO2 + 4 HF → UF4 + 2 H2 O
UF4 + F2 → UF6
- use:
separation of 235 U and 238 U in nuclear technology
Sulfur Hexafluoride (SF6 )
- manufacture: S + 3 F2 → SF6
- use:
protective gas in high voltage installations
Boron Trifluoride (BF 3 ) and Tetrafluoroboron acid (HBF 4 )
manufacture: (1)
Na2 B4 O7 + 6 CaF2 + 7 SO 3
-
use:
→ 4 BF3 + 6 CaSO4 + Na2 SO4
(reaction is carried out in conc. H2 SO4 )
(2)
HBO3 + 3 HF → BF3 + 3 H2 O
HBO3 + 4 HF → HBF 4 + 3 H2 O
(reactions are carried out in conc. H2 SO4 )
Friedel-Crafts catalyst in organic chemistry (BF 3 )
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 + COCl 2
→
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 (SiCl 4, H2, O2)
2 H 2 + O2
→ H2O
SiCl 4 + 2 H2 O
→ SiO2 + 2 HCl
(2)
direct formation from the elements in a flame of 2000 °C
(Daniell burner - 8 % of world market)
H2 + Cl 2
→
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
→
Na2 SO4 + HCl
Manufacturing in the laboratory:
2 NaCl + H2SO4
→
2 Na2 SO4 + 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
Manufacturing of (1)
from the elements
hydrogen
H2 + Br 2 → 2 HBr
halogenide:
(350 °C, Pt catalyst)
(2)
Byproduct of organic
bromine substitution
reactions
Manufacturing of MOH + HBr → MBr + H2O
halogenides:
Industrial
NaBr, use in oil industry
application:
CaBr 2,
ZnBr 2
LiBr
(1)
(2)
MOH + HI → MI + H2O
or directly from the elements
TiI4 catalysts
NaI,
KI
drying agent for air
AgI
KBr
NH4Br
photography
HI
from the elements
H2 + I2 → 2 HI
(500 °C, Pt catalyst)
hydrazine + iodine
N2H4 + I2 → 4 HI + N2
pharmaceutical purposes
photography
induction of rain
Interhalogen Compounds
Electronegative partner
(valency = -1)
Electropositive
Valency
partner
Cl
+1
+3
+5
+7
Br
+1
+3
+5
+7
I
+1
+3
+5
+7
F
Cl
Br
ClF
ClF3
ClF5
BrF
BrF3
BrF5
IF
IF3
IF5
IF7
BrCl
ICl
(ICl 3)2
-
IBr
-
AB5
AB7
Structure:
AB3
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
-1
+1
+2
+3
+4
F
OF2
Oxygendifluorid1 ,
(F-O-O-F)
Dioxygendifluorid 1
-
+5
+6
-
Cl
-
Br
-
I
-
Cl2 O
(ClO, Cl2 O2 )
(Cl2 O3 )
ClO 2
(Cl2 O4 )
(ClO 3 ), Cl2 O6
(Br2 O)
(Br2 O3 )
-
I4 O9 (+3 and +4)
I2 O4
(Br2 O5 )
-
I2 O5
I2 O6
(+5 and +7)
I2 O7
+7
Cl2 O7
NOTE: Compounds should be named as oxygen fluorides, NOT as oxides!
Grey Fields – technical importance, () – not stable under standard conditions
1
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 Cl 2 + 3 HgO → HgCl 2 · HgO + Cl 2O
- use for synthesis of hypochlorites and chlorine isocyanates
- leaching agent for textiles and wood
Br2O: - 2 Br 2 + 3 HgO → HgBr 2 · HgO + Br 2O (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 + H2 SO4 (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 + Cl 2 + 2 H2O + NaCl
- 2 NaClO2 + Cl 2 → 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 Cl 2)
I2O5: - Thermal treatment of iodine acid (200 °C)
2 HIO3 → I2 O5 + H2O
Halogen Acids and Their Salts – Overview
Nomenclature:
HXO – hypohalogenous acid
HXO2 – halogenous acid
HXO3 – halogenic acid
HXO4 – perhalogenic acid
XO- - hypohalogenite
XO2- - halogenite
XO3- - halogenate
XO4- - perhalogenate
Acids:
→ protons are bonded with an oxygen atom
Increase of acid strength
Valency
F
-1
(HOF)
+1
+2
+3
-
+4
+5
-
+6
+7
-
Cl
(HClO)
Ks = 2.9 * 108
(HClO 2 )
Ks = 1.1 * 102
(HClO 3 )
Ks = 5.0* 102
HClO 4
Ks = 1010
Br
(HBrO)
Ks = 2.1 * 10-8
(HBrO 2 )
I
(HIO)
Ks = 2.3 * 10-11
(HIO 2 )
(HBrO 3 )
Ks ~ 1
(HBrO 4 )
HIO 3
(HIO 4 )
(H5 IO 6 )
(H7 I3 O14 )
Grey fields – technical importance, () – only stable in aqueous dilution,
H5 IO6 – ortho-periodic acid, H7 I3 O14 tri-periodic acid
Base Anions:
Increase of basic strength
Valency
F
-1
(OF-)
+1
+3
+5
+7
-
Cl
ClO ClO 2 ClO 3 ClO 4 -
Br
(BrO -)
BrO 2 BrO 3 BrO 4 -
I
(IO -)
(IO 2 -)
IO 3 IO 4 H5-n IO6nH7-n I3 O14n-
Hypohalogeneous acids (HOX) and Their Salts (XO-)
HOF:
- formation at –40 °C: F2 + H2 O → HOF + HF
- decomposition in the gas phase and in weak basic solutions:
2 HOF → 2 HF + O2
- decomposition in neutral and acid solutions:
HOF + H2 O → HF + H2 O2
HOCl/OCl- :
- formation reactions:
(1)
in water: Cl2 + H2 O HCl + HClO (K << 1)
(2)
2 Cl2 + 3 HgO + H2 O → HgCl2 · HgO + 2 HOCl
(3)
in basic solutions: Cl2 + 2 OH- → Cl- + OCl- + H2 O
(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 H2 O
→ Ca(OCl) 2 · NaOCl · NaCl · 12 H2 O
+ Ca(OCl)Cl → Ca(OCl) 2 2 H2 O
+ 2 NaCl + 10 H2 O
(5)
PPG process: Ca(OH)2 + 2 HOCl→ Ca(OCl)2 +2 H2 O
(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” ([Na3 PO4 · 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 ε 0 (HClO/Cl-) = + 1.49 V
- very weak acid (K s = 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- + BrO 3 - only alkali salts are stable until 0 °C
HOI/OI- :
- formation reactions:
(1)
2 I2 + 3 HgO + H2 O → HgI2 · HgO + 2 HOI
(2)
in basic solutions: I2 + 2 NaOH → NaI + NaOI + H2 O
- very poor stability of acid, poor stability of salts
- disproportionation: 5 HIO → HIO 3 + 2 I2 + 2 H2 O
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 + H2 O2
→ 2 MClO2 + O2 + H 2O
- 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)
HIO2/IO2-:
- both very unstable, no chemistry is known
disproportionation of BrO2 BrO- → Br- + BrO2(2)
comproportionation
(solid reaction in absence of
water) Br - + BrO3- → 2 BrO2- decomposition in acid solutions, forming bromine
Halogenic acids (HOX 3) and Their Salts (XO3-)
HClO 3 /ClO 3 -:
(1) 2 HClO + ClO - → ClO 3 - + 2 HCl
(disproportionation of HClO
in acid solutions)
(2) 3 Cl2 + 6 OH- → ClO 3 - + 5 Cl- + 3 H2 O
(3) Electrolysis of NaCl without cell separation
NaCl + 3 H2 O → NaClO 3 + 3 H2
(technical process)
- acid is stable up to 40 %
- very strong oxidation agent in acid solution
e.g. ClO 3 - + 5 X- + 6 H+ → XCl + 2 X2 + 3 H2 O
- 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
- formation:
HBrO 3 /BrO 3 -: - formation: (1) 3 Br2 + 6 OH- → BrO 3 - + 5 Br- + 3 H2 O
(industrial process)
(2) Electrolysis of NaBr without cell separation
NaBr + 3 H2 O → NaBrO 3 + 3 H2
(technical process)
(3) Oxidation with chlorine
Br- + 3 Cl2 + 6 OH- → BrO 3 - + 6 Cl- +3 H2 O
- acid stable up to 50 %, than decomposes to Br2 , O2 and H2 O
- high oxidation power of the acid and the salts
- used for redox titrations (colourless → red-brown, in acids)
BrO 3 - + 5 Br- + 6 H+ → 3 Br2 + 3 H2 O
-industrial application in flour treatment
and in hair-setting lotions
HIO 3 /IO 3 -:
- formation:
(1) electrochemical or chemical oxidation of I2
e.g. I2 + 6 H2 O + 5 Cl2 → 2 HIO 3 + 10 HCl
(2) MClO 3 + I2 → Cl2 + 2 MIO 3
(in hot HNO3 , M = Na, K)
(3) 3 I2 + 6 OH- → IO 3 - + 5 I- + 3 H2 O
(4) NaIO 3 + H2 SO4
HIO 3 + 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 (XO 4-)
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- + H2 O → 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 + H 2O → 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- + Cl 2 + H2O → IO4- + 2 HCl
(2) Thermal disproportionation of iodates
Ba(IO3)2 → Ba5(IO6)2 + 4 I2 + 9 O2
- existence of periodate acid HIO4, H5 IO6 – ortho-periodic
acid and H7 I3O14 tri-periodic acid
(in water solution only H5 IO6)
- anions in solutions: H4IO6-, H3 IO62-, H2IO63-, IO4-, H2 I2O104- anions in salts of HIO4, H3 IO5 (meso-acid), H5 IO6, H6I2 O10,
H4I2O9 (di-periodic acids), H7I3 O14
Literature/References for Figures
(1)
Arnold Frederik Holleman, Egon Wiberg,
Lehrbuch der anorganischen Chemie,
101st edition, Berlin [u.a.] : de Gruyter, 1995
(6)
Gisbert Großmann, Jürgen Fabian,
Lehrwerk Chemie, Lehrbuch 2 „Struktur und Bindung – Aggregierte Systeme und
Stoffsystematik“,
5th edition, VEB Deutscher Verlag für Grundstoffindustrie, 1987
(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
cyanic acid – HCN
cyanide – CN-
N3
nitrogen hydrogen acid - HN3
azide – N3-
OCN
cyan acid - HNCO
isocyan acid - HOCN
SCN
thiocyan acid – HSCN
isothicyan acid - HNCS
cyanate – NCOfulminate – OCNthiocyanate – NCSisothiocyanate – 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)Cl 3, (NC)(NCS))
com- and disproportionation reactions
(e.g. (CN) 2 + 2 OH→
CN- + OCN- + H2O)
Course on Inorganic Chemistry
Chapter 6
Chalkogens
(Oxygen Group)
Overview About the Group
Group
Members
Atom Number
Rel. Atomic
Mass
Discovery
Oxygen
(O)
8
15.999
Sulphur
(S)
16
32.06
Selenium
(Se)
34
78.96
Tellurium
(Te)
52
127.60
Polonium
(Po)
84
[209.98]
1772
Scheele,
1774
Pristley
1818
Berzelius
1782
von
Reichenstein
1898
M. Curie
5 * 10-6
1 * 10-6
Percentage on
earth [Mass-%]
48.9
discoverer
unknown
(known
since
antiquity)
0.030
melting point
[°C]
boiling point
[°C]
state at room
temperature (25
°C)
and 1 bar
-218.75
119.6
220.5
449.5
2 * 10-14
(radioactive)
254
-182.97
444.6
684.8
1390
962
Electron
negativity
valence
numbers in
compounds
Reducing/
Oxidation
Power
Metallic/ Nonmetallic
character
Acid/Basic
properties of
oxides
Stability of
valence states
-2/-1
+2
+4
+6
colourless, yellow nonodourless,
metallic
tasteless gas solid (S8 )
(O2 )
3.5
2.5
red nonmetallic
(Se 8 ) and
grey
metallic
(Se 8 ) solids
2.4
-2
-2…+6
-2…+6
silver
metallic
solid
silver
metallic
solid
2.1
2.0
-2…+6
-2…+6
oxidising
power
reducing
power
non-metallic
metallic
acid
base
Electron configuration:
s²p 4(d 10) – 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 >H2 Te > H2Po
- bonding energy H-X decreases (463 kJ/mol (H2O), 348 kJ/mol (H2S),
276 kJ/mol (H2Se), 239 kJ/mol (H2 Te))
- acid strength increases (pK s(25 °C =
15.74 (H2O), 6.92 (H2S),
3.77 (H2Se), 2.64 (H2 Te))
2- anions X 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, Hal 2O, Hal 2O3, HalO2, Hal 2O5, HalO3, Hal 2O7 (Hal = Cl, Br, I)
sulphur and higher elements:
formed in the compositions XnHal 2 (n = 2, 3, 4), XHal 2, XHal 4 and XHal 6 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
oxidation
state
S
+2
+3
XO2
X2 O5
+4
+5
SO2 –
colourless gas
Se
Te
Po
X2 O3
PoO – black
solid
(SeO 2 )n –
Se2 O5
white needles,
oxidising
agent
TeO 2 –
Te2 O5
yellow solid
PoO2 –
yellow-red
solid
XO3
+6
SO3 –
colourless
liquid
SeO 3 –
colourless
solid
TeO 3 –
yellow solid
PoO3 –
only observed
in traces
Acids/Bases
+2:
-
Po(OH)2 - basic
+4: H2 XO3
H2 SO3 –
pKs1/2 = 1.81/6.99,
moderate reducing
power
H2 SeO 3 –
strong acid
pKs1/2 = 2.62/8.3
H2 TeO 3 –
amphoteric
pKs1/2 = 2.48/7.7,
pKB1 = 2.7,
low stability
H2 PoO3 - amphoteric
+6: H2 XO4
H2 SO4 –
pKs1/2 = -3/1.89,
moderate oxidation
power
H2 SeO 4 –
strong acid
pKs2 = 1.74
o-H6 TeO 6 –
pKs1/2 = 7.7/10.95,
high oxidation
power
-
oxidation
power
acid strenght
Oxygen
Natural sources
elementary:
in compounds:
Manufacturing
in industry:
-
in laboratory:
main component of air (20.5 %)
water (88 %),
oxides and oxygen containing salts (e.g. SO42-, CO32-),
essential part of biosphere
air rectification (Linde process),
electrolysis of water
thermal or catalytic decomposition of peroxides
Oxygen species
neutral molecules:
O2, O3 (ozone)
2negative charged species: O (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 TiCl 4
medicine
fuel cells
Ozone
Natural sources
traces in atmosphere
Manufacturing
3 O2 + 285.6 kJ/mol →
in industry/laboratory:
O3
water cooling
2 O3
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 + H2 O → 2 HF + O, O + O2 → O3
ozonisator by SIEMENS
O2
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)
UV radiation
emissions
hydrocarbons, CO
emissions
Ozone in the Stratosphere (“Good Ozone”)
Stratosphere = atmospheric layer between 10 and 50 km height
Chapman cycle
-
ozone formation:
-
ozone decomposition:
O2 + hν (< 242 nm)
→
O + O2 + M (inert molecule)
O3 + hν (< 1200 nm)
O3 + O
Ο+Ο
→
Ο3
→
Ο2 + Ο
→
2 Ο2
Catalytic ozone decomposition
O3 + X
OX+ O
→
→
O2 + OX
X + O2
X
O3 + O

→
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 H2 O
→
2 H2O2 + 2 H2 SO4
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. Ag2 O → Ag)
Technical application of H2 O2
-
leaching agent
production of perborates for detergents: NaBO2 + H2O2 → NaBO3 + H2 O
-
Na2O2:
-
BaO2:
Alkali metal peroxides
production by 2 step oxidation of Na,
strong oxidation agent, use in paper and textile leaching,
use in respirators for CO2 removal
Na2O2 + CO2 → Na2 CO3 + 0.5 O2
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
300
250
200
150
100
50
0
19601970
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 H2 O – gypsum, CaSO4 – anhydrite, MgSO4 · 7 H2O,
MgSO4 · H2 O, BaSO4, SrSO4, Na2 SO4 · 10 H2O)
Industrial production of elementary sulphur
Mining
Extraction with superheated water and air under high pressure
(Frasch process)
sulphur
pressured
air
steam
steam
molten
sulphur
sulphur containing
limestone
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 P 2S5
vulcanisation of rubber
pharmaceuticals, exterminators
concrete and road building, paints
gunpowder and fireworks
Sulphur (2)
Sulphur modifications
melting
point
solid,
light
yellow
solid,
near
colourless
boiling
point
119 °C – 159 °C
liquid,
light yellow,
low viscosity
159 °C - 243 °C
liquid,
dark red-brown,
high viscosity
243 °C – 445 °C
liquid,
dark red-brown,
low viscosity
red
gaseous
blue
violet
Chain length in solid and liquid state:
α-S and β-S
S8 molecules
λ-S
π-S
µ-S
S8 molecules
Sn molecules (n = 5…30)
polymerised molecules
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 – S2O32-)
- 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 H2 S
- 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 → FeCl 2 + 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+ + HS2 H+ + S2-,
moderate reducing agent
existence of hydrogen polysulphides H2Sn, and their metal salts
Industrial important metal sulphides
- Na2 S, NaHS: prodced from Na2 SO4 + 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 (Sn O, 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:
- Properties:
- Use:
(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, submergedburner 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)
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
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 V2 O5
(V2O5 + SO2 → V2 O4 + SO3, V2O4 + 0.5 O2 → V2O5)
multi-stage fixed bed reactors
1
SO3 yield
SO3 decomposition
catalyst
1st tray
(60 % conversion)
catalyst
2nd tray
(90 % conversion)
catalyst
3rd tray
(95 % conversion)
catalyst
4th tray
(98 % conversion)
1st heat
exchanger
2nd heat
exchanger
SO3 formation
3rd heat
exchanger
temperature [°C]
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
thiosulphate
dithionite
sulphite
disulphite
sulphate
dithionate
peroxo sulphate
disulphate
peroxo disulphate
- 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 - → S2 O62condensation:
2 HSO3- → S2O62- + H2O,
2 HSO4- → S2O72- + H2O
oxidation:
2 SO32- → S2O62- + 2 e -, 2 SO42- → S2 O82- + 2 e economically important acids:
sulphurous acid H2 SO3, sulphuric acid H2 SO4
Sulphurous and Sulphuric Acid
Sulphurous acid H2SO3
Manufacture: SO2 + H2O
[H2 SO3 ]
H+ + HSO3 (K < 10-9 )
Properties:
acid is stable only in dilution, salts are stable,
moderate acid (K S1 = 1.54 · 10-2 , KS2 = 1.02 · 10-9 )
reducing agent (forming H2 SO4 / SO42-),
can be oxidised by strong reducing agents
e.g. 6 H+ + 6 e- + SO2 → H2 S + 2 H2O (Zn/HCl),
2 SO2 + 4 H+ + 4 e- → S + 2 H2O (Fe2+),
2 SO2 + 2 H+ + 4 e- → S2 O32- + H2O (HCOO-, S),
Sulphuric acid H2 SO4
Manufacture: contact process (see SO3 )
carried out in industry as a process unit
outgoing from S, H2 S or sulphidic ores
(I)
Oxidation of the starting material to SO2
(II)
Oxidation of SO2 to SO3
(III) Formation of sulphuric acid in conc. H2 SO4
SO3 + H2 SO4 → H2 S2 O7
H2 S2 O7 + H2 O → 2 H2 SO4
Properties:
very strong, oxidising acid
(K S1 = 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 , H2 S →
S)
weak reducing agent (forming peroxo disulphuric acid)
oleum = solution of SO3 in H2 SO4
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 Na2 S2O5
Manufacture: reacting NaOH + SO2 in a saturated NaHSO3 solution
Application:
photographic industry, paper industry, textile industry,
food industry, water treatment
Sodium disulphite Na2 SO3
Manufacture: reacting NaOH + SO2 in a saturated Na2 SO3 solution
Application:
photographic industry, paper industry, textile industry,
food industry, water treatment
Calcium hydrogen sulphite Na2 S2O5
Manufacture: reacting limestone + SO2
Application:
production of sulphite cellulose
Sodium thiosulphate Na2S2O3 and ammonium thiosulphate (NH4)S2O3
Manufacture: (I)
Application:
2 NaOH + SO2 + S → Na2 S2O3 + H2 O;
Na2 SO3 + S → Na2S2O3 (50-100 °C)
(II)
2 Na2 S + Na2 CO3 + 4 SO2 → 3 Na2 S2 O3 +CO2
(III) 2 NH3 + SO2 + H2 O → (NH4 )2 SO3 ,
(NH4 )2 SO3 + S → (NH4 )2 S2O3 (80-110 °C)
fixing salts in photography (formation of [Ag(S2 O3 )]and [Ag(S2 O3 )2 ]3- complexes soluble in H2 O),
anti-chlorination agent in bleaching plants
and paper industry (Cl2 → Cl-),
flue gas desulphurisation
Sodium dithionate Na2 S2O4
Manufacture:
(I)
Application:
Zinc dust process (40 °C)
Zn + 2 SO2 → ZnS2 O4,
Zn2 S2O4 +2 NaOH → Zn(OH) 2 +Na2 S2O4
(II)
Formate process
(HCOO) - + OH- +2 SO2 → S2O42- +CO2 + H2O
(III) Amalgam process
(IV) Sodium tetrahydroborate process
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 → S2 Cl2 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 SCl 2
Manufacture: S2 Cl2 + Cl2 → 2 SCl2 at low temperatures, catalyst: I2
Application:
starting material for SOCl2 production,
sulphidising and chlorination reactions
Thionyl chloride SOCl 2
Manufacture: (I)
Application:
Sulphuryl chloride SO2Cl2
reaction of SO2 or SO3 with Cl2 , SCl2 and S2 Cl2
over an activated carbon catalyst
(II)
SO2Cl2 + PCl3 → SOCl2 + POCl3
chlorination agent in organic chemistry
(producing of herbicides, pesticides, pharmaceuticals,
dyes and pigments),
non-aqueous electrolyte in galvanic cells
Manufacture: SO2 + Cl2 → SO2 Cl2 , 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 → HSO3 Cl in HSO3 Cl
Application:
mild sulphonating and chlorosulphonating agent
in organic chemistry
Fluorosulphonic acid HSO3F
Manufacture: SO3 + HF → HSO3 F in HSO3 F
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
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
Manufacture: (I)
Selenium and Tellurium
Selenium
Manufacturing:
Properties:
Application:
1.
Oxidation of anode sludge from Cu electrolysis
Ag2 Se + O2 + Na2 CO3
→ Na2 SeO 3 + 2 Ag + CO2
2.
Acidification with H2 SO4
(→ separation of SeO 3 2- and non-soluble TeO 2 )
3.
Reduction of with SO2
H2 SeO 3 + 2 SO 2 + H2O → Se + 2 H2 SO4
2 modifications in solid state
- non- metallic red selenium Se8
- semi- metallic grey selenium Se8
electronics (e.g. rectifier and photo cells,
photocopiers
Tellurium
Manufacturing:
Properties:
Application:
1.
Oxidation of anode sludge from Cu electrolysis
Ag2 Te + O2 + Na2CO3
→ Na2 TeO 3 + 2 Ag + CO2
2.
Acidification with H2 SO4
(→ precipitation of TeO 2 )
3.
Resolving of TeO 2 in base solutions
4.
Chemical reduction of with SO2
TeO 3 2- + 2 SO 2 + H2 O → Te + 2 SO42or electrochemical reduction
silver-white colour with metallic brilliance,
semiconductor
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
(ρ water = 0.9999 g/cm³, ρice = 0.9168 g/cm³)
- maximum de nsity 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
pressure
liquid water
ice
water vapour
temperature
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
Course on Inorganic Chemistry
Chapter 7
Pnictogens
(Nitrogen Group)
Overview About the Group
Group
Members
Atom Number
Rel. Atomic
Mass
Discovery
Percentage on
earth
[Mass-%]
melting point
[°C]
boiling point
[°C]
Nitrogen
(N)
7
14.007
Phosphorus
(P)
15
30.97
Arsenic
(As)
33
74.92
Antimony
(Sb)
51
121.75
Bismuth
(Bi)
83
208.98
1772
Scheele
0.017
1669
Brand
0.10
ca. 1250
Magnus
1.7 * 10-4
1492
Valentin
2 * 10-5
known since
antique
2 * 10-5
-209.99
44.25 (P 4 )
-
630.7
271.3
-195.82
280.5 (P 4 )
616
(sublimation)
yellow nonmetallic
solid (As4 ),
black nonmetallic
solid (As8 ),
grey
metallic
solid (As8 ),
1635
1580
state at room
colourless,
temperature (25 odourless,
°C)
tasteless gas
and 1 bar
Electron
negativity
valence
numbers in
compounds
Reducing/
Oxidation
Power
Metallic/ Nonmetallic
character
Acid/Basic
properties of
oxides
Stability of
valence states
–3
+3
+5
3.0
-3…+5
white nonmetallic
solid (P 4 ),
red nonmetallic
solid (P 8 ),
black semimetallic
solid (P 8 ),
purple nonmetallic
solid (P 8 )
2.1
-3…+5
grey
metallic
solid (Sb8 )
silver
metallic
solid (Bi8 )
2.0
1.9
1.9
-3…+5
-3…+5
+3 (+5)
oxidising
power
reducing
power
non-metallic
metallic
acid
base
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 (∆ B H = – 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, P 3H5)
Halogen compounds (pnictogen halogenides)
-
-
nitrogen: formed in the compositions NX3, N2X4, N2X2 and N3X as derivates from
the hydrogen compound
higher elements: XHal 3, X2Hal 4 and XHal 5 with X as the electropositive partner
formation of oxyhalogenides (e.g. NOCl, POCl 3)
General Properties of Pnictogens (2)
Oxygen compounds (pnictogen oxides)
-
-
nitrogen: formed in the compositions N2O, NO, N2 O2, N2 O3, NO2, N2O4, N2O5, NO3
and N2O6
higher elements: X2O3 and X2O5 with X as the electropositive partner
Acids/Bases
-
-
-
nitrogen: H3 NO, HNO, H2N2O2, H2N2O3, HNO2, HNO3, NO43- and HNO4
phosphorus: H3PO2, H4P 2O4, H3PO3, H2P 4O5, H2P 4O6, H3PO4, H4P 2O7 H4P 2O8 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, Bi 2O3 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. NH2 SO3H – 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:
in compounds:
Manufacturing
in industry:
in laboratory:
main component of air (79.5 %)
water (solved nitrate salts – 1 % of total N),
salt deposits (Chile, India)
essential part of biosphere (amino acids) 0.001 % of total N
air rectification (Linde process),
thermal or catalytic remo val 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-,
pK B = 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 Cl 2): 2 NH3 + ClO- → N2H4 + H 2O + 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 · H2 O (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)+ +OHApplication
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 H2 SO3 + H2O → NH2OH + 2 H2 SO4,
carried out with NH4NO2 and SO2 in diluted H2 SO4
(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 H2 O,
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 + H 2) and dinitrogen
monoxide: NaNH2 + N2 O → 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 H2 SO4
and water removal with CaCl 2
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-, pK S = 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
rectification
steam
reforming
partial
oxidation
water
nitrogen oxygen
rough synthesis gas (N2, O2, CO, H2, H2S)
absorption of CO and H 2S
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)
coal
gasification
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/FeCrO x catalyst (rest: 1 g CO /m³ air)
350 - 380 °C/sulphur insensitive Co/Mo oxide catalyst (rest: 0.3 % CO)
Removal of CO and H2 S
- 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 /N 2 = 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 % Fe3 O4 , 0.8 % K2O, 2.3 % Al2 O3 , 1.7 % CaO, 0.5 % MgO,
0.4 % SiO 2
preparation:
(1)
melting a mixture of magnetite (Fe3 O4 ) 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
(Fe3 O4 + 4 H2 → 3 Fe + 4 H2 O)
Fe - active component
K2O acts as electronic promoter (increase of activity)
Al2 O3 and SiO 2 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
pressure
resistent
outer tube
catalytic
contact
heat
exchanger
tube
temperature
tubular reactor
multiple-bed reactor
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
Oxides
Acids
+1
N2O
HNO,
H2N2O2
+2
NO,
N2O2
N2O32-
+3
N2O3
HNO2
+4
NO2,
N2O4
HNO3
NO43-
+5
N2O5
HNO4
Oxides:
endothermic meta-stable compounds (without N2O4 and N2O5)
occur during high temperature combustion processes from nitrogen oxidation
(instead of N2O)
equilibriums NO/N2 O2 and NO2/N2O4 (dimerisation primarily at higher
temperatures)
N2O, NO/N2O2 and NO2/N2 O4 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
N2O32- and NO43- exist only as salts
Nitrogen Oxides
Dinitrogen monoxide N2O (NNO)
Industrial manufacturing
careful heating of ammo nium 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 H2 O + 904 kJ/mol
(2)
further oxidation of NO to NO2 /N 2O4 at ca. 150 °C
2 NO + O2 → 2 NO2 + 114 kJ/mol, 2 NO2 → N2 O4 + 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
NO
Pt-Rh alloy gauze
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 N2 O4 (head)
(5)
conversion of N2 O4 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 H2 SO4
- 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 + H2 O + 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 H2 O)
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 H 2O + 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 + P 2 (g)
(reduction of P 2O5 by C, SiO2 is added to form a slag with Ca)
condensation and distillation of rough phosphorus
→ “white phosphorus”
carbon electrode
gas outlet (P 2)
carbon electrode
phosphatite
outlet
for slags
phosphate
carbon
electrodes
outlet
for slags
electrode mass
view from side
view from top
outlet
for molten
iron
from
electrode
mass
Elementary Phosphorus (2)
Modifications
White
Phosphorus
electro-thermal
reduction of
phosphates
Red
Phosphorus
thermal treatment
of white P
at 200-400 °C
for 20-30 hours
(∆BH =
- 17.7 kJ/mol)
Violet
Phosphorus
thermal
treatment of
white P
at > 550 °C
for 1-2 weeks
Molecular and
crystal structure
P4 tetrahedrons
P8 ,
amorphous
P8 ,
complex layer
structure
Stability
meta-stable at
25 °C,
thermodyn.
stable at
> 620 °C
44.3 °C
meta-stable at 25
°C
meta-stable at
25 °C,
thermodyn.
stable
at 550-620 °C
sublimation at
620 °C
low
low,
flammable at
> 400 °C
Manufacture
Melting point
Boiling point
Toxicity
Reactivity
Metallic/
non- metallic
character,
electrical
conductivity
280.5 °C
high (T+)
extremely high,
high flammable,
strong reducing
agent
non- metallic
isolator
sublimation at ~
580 °C
low
moderate,
heavy reaction
only with strong
oxidation agents
non- metallic
isolator
non- metallic
isolator
Phosphorous species in compounds
- occur in all oxidation states from –3 to +5
-
Application
90 % for manufacture of P2 O5 (→ phosphorous acid, phosphates)
synthesis of P-S and P-Halogen compounds (→ organic chemistry)
safety matches (red phosphorous)
military purposes
Black
Phosphorus
thermal treatment
of white P
at 380 °C for
some days
(adding of
dispersed Hg as
catalyst)
P8 ,
layer structure
(similar to
graphite)
stable form
at < 550 °C
changes to violet
P
low
low,
flammable at
> 400 °C
metallic
semiconductor
Phosphorus Oxides
Phosphorus trioxide (P4O6) and phosphorus pentoxide (P4O10)
phosphorus trioxide (P4O6)
-
-
-
manufacture:
phosphorus pentoxide (P4O10)
combustion of P 4 at low temperatures and oxygen
shortage (P4O6) or with dried air under O2 excess
(P4O10)
P 4 + 3 O2 → P 4O6 + 1641.2 kJ/mol
P 4 + 5 O2 → P 4O10 + 2786 kJ/mol,
separation of oxides by fractioned vaporisation
properties:
P 4O6 – white wax-like high toxic solid,
melting point: 23.8 °C,
boiling point: 175.3 °C (under N2),
oxidation to P 4O10 at temperatures > 70 °C
(∆RH = + 672.4 kJ/mol),
reaction with cold water to phosphonic acid (H3PO3),
with HCl to H3PO3 and PCl 3 and
with halogens to phosphoryl halogenides
P 4O10 – 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: P 4O6 has no technical importance,
P 4O10 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
acids
phosphinate
phosphonate
“+1”
“+2”
“+3”
H3PO2
H4P 2O2
(HPO2)n,
H3PO3,
H4P 2O5
phosphate
peroxophosphate
-
“+4”
+5
H4P 2O6
(HPO3)n,
H3PO4,
H3PO5,
H4P 2O7,
H4P 2O8
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 diphosphonate
(diphosphate(III)) (diphosphate(III))
hypodiphosphate
(diphosphate(IV))
diphosphate(II,IV) diphosphate(III,V)
diphosphate
(diphosphate(V))
peroxodiphosphate(V)
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 H2 O),
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 P 4 in air excess and
absorption of the formed P 4O10 together
with water in conc. H3PO4 (85 %)
P 4 + 5 O2 → P 4O10, P 4O10 + 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 → H4P 2O7 + H2 O
concentrated (85 %) acid = high viscous liquid,
melting point: 21.1 °C, boiling point: 158 °C,
three-base middle-strong acid
(pK S1 = 2.16, pKS2 = 7.21, pK S3 = 12.32),
pure acid = strong oxidation agent > 400 °C,
diluted acid = no oxidising properties
buffer area
pH
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 H 3PO2
-
-
Ortho-phosphorous/ortho-phosphoric/ortho-phosphonic acid (H3PO3)
manufacture:
(I)
PCl 3 + 3 H2O → H3PO3 + 3 HCl at 190 °C
(II) P 4O6 + 6 H2O → 4 H3PO3
properties:
colourless crystals, melting point: 73.8 °C,
middle-strong two base acid (pK S1 = 2.0, pK S2 = 6.6),
in aqueous solutions high solubility of alkali salts,
low solubility of other salts,
phosphoric acid
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, H2 SO4 to H2SO3,
stable under air atmosphere at room temperature
disproportionation at heating of the dry acid
phosphonic acid
4 H3PO3 → PH3 + 3 H3PO4 (130 – 140 °C)
Organyl derivates are known
from both tautomeric forms .
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 P 4 with NaOH or Ca(OH) 2
2 P 4+ 3 Ca(OH) 2 + 6 H2O
→ 2 PH3 + 3 Ca(H2PO2)2↓
hypophosphoric acid
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 P 4 with warm water
(disproportionation)
phosphinic acid
P 4 + 6 H2O → PH3 + 3 H 3PO2
properties:
colourless flakes, melting point: 26.5 °C,
middle-strong one base acid (pK S = 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)
Organyl derivates are known
from both tautomeric forms .
Salts and Organic Compounds from Phosphorus Acids (1)
Fertilizers - general importance
- 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 → Na2H2P 2O7 + 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 % P 2O5,
double superphosphate
→
35 % P 2O5,
triple superphosphate
→
> 46 % P 2O5
ammonium phosphates (made by neutralisation and
thermal condensation), used purely and
in H2 O solution (high solubility)
“Rhenania phosphates” – made by sintering of apatite
with silica and Na2CO3 or NaOH (29 % P 2O5,
low solubility)
“Melt phosphates” – made by melting apatite with
Mg compounds and silica (21 % P 2O5, low solubility)
“Thomas phosphates” = slag from smelting
P containing iron ores (10-18 % P 2O5, 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
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
in animal
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, P 2X4, PX5, POX3, PSX3 (X = F, Cl, Br, I), P 4F6, P 6Cl6, P 6Br6
mixed halogen compounds and partial substitution of halogens by hydrogen are
possible
technical importance: PCl 3, PCl 5 and POCl 3
synthesis:
PCl 3 direct conversion of white phosphorus with dry chlorine
¼ P 4 + 1.5 Cl 2 → PCl 3 + 320 kJ/mol
PCl 5 chlorine addition to PCl 3, PCl 3 +Cl 2 → PCl 5 + 124 kJ/mol
POCl 3 oxidation of PCl 3 with oxygen at 50 – 60 °C
PCl 3 +1/2 O2 → POCl 3 + 277.6 kJ/mol,
synthesis from P 4O10 and PCl 3
P 4O10 + 6 PCl 3 + 6 Cl 2 → 10 POCl 3
PSCl 3 PCl 3 +S → PSCl 3 in autoclaves at 180 °C
properties:
PCl 3 colourless, smoking, toxic liquid with stabbing odour,
hydrolysis in water to H3PO4 and HCl, Lewis base properties,
moderate reducing agent
PCl 5 green-white toxic solid, decomposes at higher temperatures
into PCl 3 and Cl 2, in the presence of water
via POCl 3 + HCl to H3PO4 and HCl,
Lewis acid
POCl 3 colourless, smoking, toxic liquid
PSCl 3 colourless liquid,
melting point: -35 °C, boiling point: 125 °C,
decomposes with water to H3PO4, HCl and H2S
use:
PCl 3 synthesis of H3PO3 (10 %) and alkyl substituted phosphates
(detergents), of PCl 5, POCl 3 (33 %) and PSCl 3,
starting material for ligand compounds
in metal-organic chemistry
PCl 5 chlorination agent in organic chemistry
POCl 3 manufacture of POX derivates (X = -OR, -NHR, -R)
used as lubricant additives, softeners,
flame inhibitors and insecticides
PSCl 3 manufacture of thiophosphoric acid ester chlorides
for the production of pesticides
Other Industrial Important Phosphorus Compounds (2)
Phosphorus hydrogen compounds (phosphanes)
PH3, P 2H4, P 3H5 (linear compounds), P 5H5, P 7H3 (cyclic compounds)
with P as the electronegative partner
PH3 – exothermic compound, all other P xHy are endothermic compounds
Technical synthesises of PH3:
(1)
P 4 + 3 NaOH + 3 H2O
→
PH3 + 3 NaH2PO2
(2)
2 P 4 + 12 H2O
→
5 PH3 + 3 H3PO4
Properties of PH3:
colourless, toxic gas with garlic odour,
low solubility in water, neutral reaction
(pK B (PH3/PH4+)= 27, pKS (PH3/PH2-) = 29),
salts: phosphides P 3- (wide variety of higher phosphides from P xHy)
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:
P 4Sn (n = 2 – 10), PSn (phosphorus polysulphides),
different thiophosphates (linear compounds),
P 4Sen (n = 3-5), P 2Se5
manufacture: fusing of P 4 and S/Se, exothermic reactions;
seperation by extraction with CS2,
com- and disproportionation reactions
application: P 4S10 has some technical importance
(flotation agent, lubricant additive,
manufacture of insecticides)
Phosphorus nitrogen compounds
polymeric (NPCl 2)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
Course on Inorganic Chemistry
Chapter 8
Carbon Group
Overview About the Group
Group Members
Atom Number
Rel. Atomic
Mass
Discovery
Percentage on
earth
[Mass-%]
density [g/cm³]
melting point
[°C]
boiling point
[°C]
state at room
temperature (25
°C)
and 1 bar
Electron
negativity
valence numbers
in compounds
Reducing/
Oxidation Power
Metallic/ Nonmetallic
character
Electrical
conductivity
Acid/Basic
properties of
oxides
Stability of
valence states
–4
+2
+4
Physiology
Carbon
(C)
6
12.01
Silicon
(Si)
14
28.09
Germanium
(Ge)
32
72.59
Tin
(Sn)
50
118.69
Lead
(Pb)
82
207.2
unknown
1824 Berzelius
26.3
known since
antique
2 * 10-4
unknown
0.02
1886
Winkler
1.4 * 10-4
3.51
(diamond)
2.26
(graphite)
2250
(sublimation)
2.32
5.32
7.28
11.34
1410
937.4
231.9
327.4
2477
2830
2687
1751
colourless
diamond ,
polymeric
black graphite,
fullarenes
2.5
-4…+4
dark-grey,
hard, brittle
metal
grey-white,
very brittle
metal
silver-white,
very soft
metal
1.2 * 10-5
blue-grey,
very soft
metal
1.8
1.8
1.8
1.8
-4…+4
-4…+4
-4…+4
+4 (+2)
oxidising
power
non-metallic
(diamond),
semimetallic
semimetallic
semimetallic
(graphite)
isolator
(diamond),
semi -conductor semi -conductor
semi -conductor
(graphite)
reducing
power
metallic
metallic
conductor
conductor
acid
acid
amphoteric
amphoteric
primarily
base
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 (∆ B H = – 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: XHal 2 and XHal 4 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:
-
in compounds:
diamonds (Africa, Brazil, Siberia)
graphite (Madagascar, Sri Lanka, Korea, Norway etc.)
carbonates (lithosphere, hydrosphere),
organic compounds (biosphere);
carbon dioxide (atmosphere – 0.03 %, hydrosphere)
Modifications
elt
/m
nd
mo
dia
diamond
pressure
gra
ph
ite
/m
elt
(m
eta
-
diamond
ite /dia
graph
liquid
carbon
sta
ble
)
mond
graphite
gaseous
carbon
temperature
diamond
-
diamond:
-
graphite:
-
graphite
p-T diagram of carbon
- colourless non-metallic modification,
- electrical isolator
- high hardness (used for tools)
- manufacture: mining or treating of graphite
at high temperatures and extreme pressures
- 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 (∆R H = + 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 C4 H10
straight and branched aliphatic compounds with multiple π−π C-C bonds
(unsaturated hydrocarbons - olefins)
ethylene C 2H4
cyclic compounds
cyclohexane
propylene C3H6
butadiene C4H6
acetylene C2H2
propine C3H4
aromatic compounds
cyclopentadiene
benzene
naphthalene
anthracene
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
temperature
-
conversion of coal is performed at 1000 °C in Winckler generators
product mixture = “generator gas” (70 % N2, 25 % CO and 4 % CO2)
Industrial Carbon Oxide Chemistry (2)
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 C2 H2n+2 + n H2 O → n CO + 2(n+1) H2
(2)
Partial Oxidation of heavy fuel oil
2 C2 H2n+2 + n O2 → 2n CO + 2(n+1) H2
(3)
Coal gasification
C + H2O
CO + H2, 2 C + O2
2 CO,
Industrial Carbon Oxide Chemistry (3)
Synthesis Gas
-
Application:
ammonia
urea
resins
formaldehyde
polyols
synthesis
oxidation
paraffins
coal
acetates
olefins
alcohols
acetic acid
anhydride (ESA)
synthesis gas
(syngas)
methanol
synthesis
ethylene
aromatic comp.
min. oil/
nat. gas
polypolymethylene
merisation
Mobile process
olefins
methanisation
process
+ isobutene
oxosynthesis
oxo-aldehydes
oxo-alcohols
fermentation
fuels
methyl-tertbuthyl-ether (MTBE)
Further Important Carbon Compounds
-
-
-
Carbides
compounds with electropositive elements (anions C4-, C22-)
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 MC 2 :
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 30004000 °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. SnO 2 → 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, ∆ R H = 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 SiHCl 3 (300 °C)
Si + 3 HCl → SiHCl 3 + H2
(2)
Distillation of SiHCl 3
(3)
decomposition of SiHCl 3 at 1000 °C
yielding to highly purified Si
SiHCl 3 + 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
cyclic compounds (Si > 4)
cyclopentasilane
→
→
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 SiH 4 :
manufacture: (1)
-
properties:
-
application:
decomposition of Mg2 Si with acids
in the absence of air
Mg2 Si + 4 H+ → SiH4 + 2 Mg2+
(2)
hydrogenation of SiCl4 with LiH
in molten LiCl/KCl at 400 °C
SiCl4 + 4 H- → SiH4 + 4 Cl- 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 → SiO 2 + 2 H2 O + 1518 kJ/mol
- in the presence of water:
SiH4 + 2 H2 O → SiO 2 + 4 H2 + 374 kJ/mol
- with halogens and hydrogen halogenides
stepwise replacing of H by Hal
manufacture of ultra pure Si
n
polysilane
H3Si-(SiH2 )n-SiH 3
Technical Important Binary Silicon Compounds
-
-
-
Silicon halides:
SiF 4 manufacture: CaF2 + H2 SO4 → 2 HF + CaSO 4
SiO 2 + 4 HF → SiF 4 ↑ + 2 H2 O (in conc. H2 SO4 )
properties:
- highly exothermic compound, stable under dry air,
decomposes in water:
3 SiF 4 + 2 H2 O → SiO 2 (aq) + 2 H2 SiF 6
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
manufacture:
-
properties:
-
application:
-
manufacture:
-
properties:
-
application:
Silicon dioxide (SiO 2 ) – quartz
- mining of quartz sands, re-crystallisation in H2 O
- flame hydrolysis of SiF 4 and SiCl4
- high stable, exothermic solid
- crystal and glass (amorphous) modification
- reacts only with HF at room temperature,
- reacts with alkali hydroxides in molten state
- 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 (Si3 N4 )
(1)
from the elements at 1100-1400 °C
3 Si + 2 N2 → Si3 N4 + 750 kJ/mol
(2)
3 SiO 2 + 2 N2 + 6 C → Si3 N4 + 6 CO
(3)
3 SiCl4 + 4 NH3 → Si3 N4 + 12 HCl
- colourless solid with high hardness
but low density
- stable up to 1900 °C
- resistant against corrosion and mechanical stress
- special ceramics
- manufacturing of chemical apparatus,
high quality mechanical tools
- mechanical and motor engineering
- fittings
Silicon carbide (SiC):
Metal Silicides:
see carbides under
“Further Important Carbon Compounds”
metallic hard materials
Silicon Acid and Silicates
Manufacture: - Silicates:
melting of alkali carbonates or hydroxides
with quartz sand
- Acid:
(1) solving alkali silicates in H2 O and
precipitation of the acid by slow acidification
(2) hydrolysis of monomeric SiX4 compounds
SiX4 + 4 H2O → “H4SiO 4” + 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
Application:
Tertiary structure of zeolithe A
- 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 CH3 Cl → (CH3)2SiCl 2
Si + 2 PhCl → Ph2-SiCl 2
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 (H2 SO4 )
-
-
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)
-
-
-
-
-
-
Manufacturing of the metal
outgoing from waste gases of Zn manufacture
Process steps:
(1)
extraction GeO2 and ZnO of fly ash with H2 SO4
(2)
precipitation of oxides at pH = 5 with NaOH
(3)
conversion of oxides to chlorides with HCl
(4)
distillation → separation of GeCl 4 and ZnCl 2
(5)
Hydrolysis of GeCl 4 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
Manufacturing of the 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
SnCl 4: - obtained from the elements by treating of tinplate wastes
with chlorine: Sn + 2 Cl 2 → SnCl 4 + 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
Manufacturing of the metal
“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,
PbCl 2, 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-charging
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
energy
Course on Inorganic Chemistry
Chapter 9
Earth Metals
(Boron Group)
Overview About the Group
Group Members
Atom Number
Rel. Atomic
Mass
Discovery
Percentage on
earth
[Mass-%]
density [g/cm³]
melting point
[°C]
boiling point
[°C]
Electron
negativity
valence numbers
in compounds
Metallic/ Nonmetallic
character
Electrical
conductivity
Acid/Basic
properties of
oxides
Stability of
valence states
+1
+3
Physiology
Boron
(B)
5
10.81
Aluminium
(Al)
13
26.98
Gallium
(Ga)
31
69.72
Indium
(In)
49
114.82
Thallium
(Th)
81
204.37
1808
Gay-Lussac,
Thenard, Davy
0.001
1825
Oerstedt
1875
de Boisbaudran
1861
Crookes
7.7
1.6 * 10-3
1863
Reich,
Richter
1 * 10-5
2.46
2.70
5.91
7.31
11.85
2250
(sublimation)
660.3
29.8
156.6
303.5
2330
2403
2070
1453
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
isolator
conductor
conductor
conductor
conductor
acid
amphoteric
amphoteric
amphoteric
base
not essential,
non-toxic
not essential,
non-toxic
-
-
toxic
Electron configuration:
s²p1 (d10 ) – need of loosing 3 electrons
for full saturation of electron shells
5 * 10-5
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
Hydrogen compounds and metal salts (-ides)
- covalent compounds
- formation of dimeric B2 H6 and of polymeric (AlH3)n
- GaH3, InH3, TlH3 were not found
- stability of hydrogen compounds:
½ B2H6 (∆ B H = +18 kJ/mol) < 1/n (AlH3)n (-11 kJ/mol)
Halogen compounds
-
B:
Al:
Ga, In, Tl:
BHal 3, B2Hal4 and (BHal)n compounds
polymeric compounds
XHal, X-XHal 4 and XHal 3 compounds
Binary compounds with chalkogenes
-
X2 Y3, Ga, In, Tl forms additionally mono-chalkogenides X2 Y
Binary compounds with pnictogenes
-
XY – hard compounds, diamond-like structure, partially semiconductors
Acids/Bases
-
B:
Al, Ga, In:
Tl:
H3BO3 – acid properties
amphoteric X(OH) 3 compounds
weak base Tl(OH) 3, strong base TlOH
Elementary Boron
Natural sources
occurs only in compounds
minerals: kernite Na2B4O7 · 4 H2O and borax Na2B4 O7 · 10 H2 O
Manufacture of the element
amorphous boron (technical grade)
thermal reduction of B2 O3 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 BHal 3 + 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 Cl 2,
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 Bn Hn+2 , BnHn+4 , BnHn+6 , BnHn+8 , BnHn+10),
as anions (e.g. BH4-, B3 H82-),
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 –
Na2B4 O7 · 7 SiO2 + 16 Na + 8 H2
→ 4 NaBH4 + 7 Na2SiO3
Halogen compounds
-
-
-
BHal 3, 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
BCl 3: - manufacture: B2 O3 + 3 C + 3 Cl 2 → 2 BCl 3 + 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 - Na2B4 O7 · 10 H2 O – used for ceramics, enamels, glassware
perborates - NaBO2 · 2 H2O – used in detergents as leaching agent
Boron Compounds (2)
Boron nitride BN
manufacture:
-
properties:
use:
Boron carbide B4C
manufacture:
-
properties:
use:
B2O3 + 2 NH3 → 2 BN + 3 H2 O (800-1200 °C) or
B2O3 + 3 C + N2 → 2 BN + 3 CO
highly inert material
high temperature lubricant,
lining of rocket burning chambers, plasma burners and
nuclear reactors
B2O3 + 7 C → B4C + 6 CO (at 2400 °C) or
2 B2O3 + C + 6 Mg(4 Al) → B4C + 6 MgO + 2 Al 2 O3
highly inert and hard material (similar than diamond)
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 2 O3
electrolysis of cryolyth (82 %)-alumina (18 %) mixture at 940-980 °C
molten electrolyte
carbon anode
isolation
liquid Al
carbon cathode
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 Fe 3O4 + 8 Al → 4 Al 2O3 + 9 Fe + 3341 kJ)
important compounds: Al 2O3, Al(OH) 3, Al 2(SO4)3, AlCl 3, NaAlO2, AlF3, NaAlF6,
spinells
Course on Inorganic Chemistry
Chapter 10
Alkaline Earth Metals
Overview About the Group
Group
Members
Atom Number
Rel. Atomic
Mass
Discovery
Beryllium
Be
4
9.01
Magnesium
Mg
12
24.31
Calcium
Ca
20
40.08
Strontium
Sr
38
87.62
Barium
Ba
56
137.33
1828
Wöhler
1809
Davy
1790
Grawford
1774
Scheele
2.0
1808
Berzelius,
Pontin
3.4
Radium
Ra
88
[226]
radioactive
1898
Curie
3.6 * 10-2
4 * 10-2
1 * 10-10
648.8
839
768
710
~ 700
1105
1482
1380
1537
~ 1140
1.74
1.54
2.63
3.65
unknown
1.2
1.0
1.0
0.9
0.9
+2
+2
+2
+2
+2
Percentage on
2.7 * 10-4
earth
[Mass-%]
melting point
1278
[°C]
boiling point
~ 2500
[°C]
Density [g/cm³]
1.85
at 25 °C
and 1 bar
Electron
1.5
negativity
valence
+2
numbers in
compounds
Reducing/
moderate
Oxidation
reducing
Power
agent
Metallic/ Nonmetallic
character
Acid/Basic
amphoteric
properties of
oxides
Physiology
highly toxic
strong
reducing
agent
silver-white or yellow-white metals
base
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
Hydrogen compounds and metal salts (-ides)
- 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 (∆B H = -184 kJ/mol)
> SrH2 (∆ B H = -177 kJ/mol) > baH2 (∆B H = -172 kJ/mol)
- salt-like hydrides (except “BeH2”), stable under air
- reaction with water to hydroxides and H2
Halogen compounds
- Be:
- higher elements:
Chalkogen compounds
- Be:
- higher elements:
Acides/Bases
- Be:
- higher elements:
covalent halogen compounds (BeHal 2)n
with Lewis acid character
ionic salts XHal 2
BeY, stable in air and water
XY, stable in air, hydrolysis in water
(forming hydroxides and H2 Y)
Be(OH) 2, soluble in acid and base solutions
X(OH) 2, only soluble in acids, base reaction
Beryllium (1)
Natural sources
- rather rare element
- minerals:
beryl - 3 BeO · Al 2O3 · 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 H2 SO4
- 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 + Cl 2 → BeCl 2 + CO at 800 °C
purification of BeCl 2 by distillation at 485 °C
electrolysis of molten, water-free BeCl 2:
BeCl 2 → Be + Cl 2 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+
soluble in
acid solutions
Be(OH) 2
precipitation
in neutral solutions
Be(OH) 42soluble 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 · MgCl 2 · 6 H2 O,
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:
(2)
thermal conversion of MgO
MgO + C + Cl 2 → MgCl 2 + CO + 153 kJ/mol,
electrolysis of molten, water-free mixture of MgCl 2 (8 - 24 %)
and other alkali and alkaline earth metal chlorides,
MgCl 2 → Mg + Cl 2 at 700 - 800 °C, removal of molten Mg
(favoured process, 80 % of world Mg)
thermal reduction of dolomite with ferrosilicon (vacuum, 1200 °C)
2 (CaO · MgO) + Si(Fe) → 2 Mg + Ca2 SiO4 (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 )2 CO3 → MgCO3 ↓ + 2 NH4 + or
by carbonating of MgO (e.g. obtained from dolomite)
MgO + 2 CO2 + H2 O → 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 H2 O)
- 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:
- Application:
Magnesium (3)
(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 + H2 O
temperature: 600-900 °C
- “caustic” MgO,
1600-2000 °C - “sintered” MgO,
melting at 2800-3000 °C
– “fused”/”dead burnt” MgO
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
Important magnesium salts (continuation)
Magnesium Chloride (MgCl2 )
- Manufacturing:
(1)
- Application:
from brines and seawater
Dow Chemical process
- precipitation of Mg(OH)2 with lime
- Mg(OH)2 + Ca(OH)2 + 2 HC l + 2 H2 SO4
→ MgCl2 + CaSO4 ↓ + 4 H2O
- evaporation at 200 °C → MgCl2 · 2 H2 O
- 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
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 H2 SO4
- Application:
fertilizer (80 %), manufacture of K2 SO4 , Na2 SO4 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 · Al 2O3 (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 SmCo 5 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
CaCO 3
(limestone)
(1)
CaO
(quicklime)
(2)
Ca(OH)2
(slake lime/
lime hydrate)
(1)
calcinations of limestone at 1000 – 1200 °C
178.4 kJ/mol + CaCO3 → CaO + CO2
(2)
slaking of CaO = slow addition of water
CaO + H2 O → Ca(OH)2 + 65.2 kJ/mol
use of slake lime in construction,
binding of slake lime with atmospheric carbon dioxide
Ca(OH)2 + CO2 → CaCO 3 + H2O
(3)
cement
=
lime mortar
gypsum mortar
=
=
(3)
CaCO 3
(limestone)
3 CaO · SiO 2 ,
obtained by thermal treating
of a mixture of CaO, SiO 2 and some additives
(Fe2 O3 , Al2 O3 ) at 1450 °C
mixture of slake lime and sand
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,
Na2 CO3 – Solvay process)
- agriculture and sugar industry
- refractory materials
- flue gas desulphurisation
Other Calcium Compounds
Calcium carbide (CaC2)
-
manufacturing:
-
application:
reacting highly purified CaO with coke
in an electrical furnace at 2000 – 2200 °C
464 kJ/mol + CaO + 3 C → CaC2 + CO
- formation of acetylene
CaC2 + 2 H2 O → 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:
application:
CaC2 + N2 → CaCN2 + C + 296 kJ/mol (700–900 °C)
- 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 (CaCl 2)
-
-
manufacturing:
waste product from many processes
- soda manufacturing (Solvay process)
CaCO3 + 2 NaCl → Na2 CO3 + CaCl 2
- 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:
- application:
Sr(NO3)2: - manufacturing:
- application:
mining
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)
SrCO3 + 2 HNO3
fireworks
→ Sr(NO3)2 + CO2 + H2O
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:
- application:
BaCO3:
- manufacturing:
- application:
BaO2:
(1)
mining
(2)
oxidation of BaS with Na2 SO4
drilling-mud for oil and gas exploration
(90 % of mined BaSO4),
white pigment (manufacture of paper, paint,
rubber and plastics)
3 step process
(1)
BaSO4 + 4 C → 4 BaS + 4 CO
(rotary kiln, 1200 °C)
(2)
BaS + CO2 + H 2O → BaCO3 ↓ + H2S or
BaS + Na2CO3 → BaCO3↓ + Na2S
(in aq. solution)
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
- manufacturing:
- application:
Ba(NO3)2: manufacturing:
- application:
glowing of BaCO3 and coke
BaCO3 + C → BaO + 2 CO,
thermal oxidation of BaO at 500-600 °C
2 BaO + O2 → 2 BaO2
igniting agent
BaCO3 + 2 HNO3 → Ba(NO3)2 + CO2 + H2O
fireworks
Course on Inorganic Chemistry
Chapter 11
Alkali Metals
Overview About the Group
Group
Members
Atom Number
Rel. Atomic
Mass
Discovery
Percentage on
earth
[Mass-%]
melting point
[°C]
boiling point
[°C]
Density [g/cm³]
at 25 °C
and 1 bar
state at 25 °C
and 1 bar
Electron
negativity
valence
numbers in
compounds
Reducing/
Oxidation
Power
Metallic/ Nonmetallic
character
Acid/Basic
properties of
oxides
Lithium
(Li)
3
6.94
Sodium
(Na)
11
22.99
Potassium
(K)
19
39.10
Rubidium
(Rb)
37
85.47
Caesium
(Cs)
55
132.91
1817/18,
Arfevedson,
Davy
2.0 · 10-3
1803,
Davy
1807,
Davy
2.7
2.4
1861/62,
Bunsen,
Kirchhoff
9.0 · 10-3
1860,
Bunsen,
Kirchhoff
3.0 · 10-4
180.5
97.8
63.6
38.9
28.4
~ 27
1347
881.3
753.8
688.0
678
~ 660
0.53
0.97
0.86
1.5
1.9
unknown
0.7
+1
1.0
0.9
0.8
0.8
malleable
golden
metal
0.7
+1
+1
+1
+1
+1
malleable silver metals
Electron configuration:
strong reducing agents
typical metals
basic oxides and hydroxides
s1 (d10 ) – need of loosing 1 electron
for full saturation of electron shells
Francium
(Fr)
87
[223]
(radioactive)
1939, Perey
1.3 · 10-21
General Properties of Alkali Metals (1)
Manufacture of metals
electroylsis of molten, water-free salts (Downs process):
2 LiCl → 2 Li + Cl 2 at 610 °C
2 NaCl → 2 Na + Cl 2 at 600 °C
chemical reduction:
KCl + Na → K + NaCl (850 °C)
2 RbOH + Mg or Ca
→ Mg(OH) 2 or Ca(OH) 2 + Rb
Cs 2Cr2O7 + 2 Zr
→ 2 Cs + 2 ZrO2 + Cr 2O3
(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
General Properties of Alkali Metals (2)
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 + Hal 2
→
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 Cl 2
(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
General Properties of Alkali Metals (3)
Chalkogen compounds (general)
-
-
formation of compounds M2 Y (chalkogenides) with partially covalent character,
M2 Yn>1 (perchalkogenides) and M>2 O (suboxids)
properties of sulphides → chapter 7
Oxides, suboxides, peroxides and ozonides
Li
Na
K
Rb
Cs
Oxides M2O
Peroxides M2O2
Hyperoxides MO2
Ozonides MO3
Li2O2
Na2O1
K2O6
Rb2O6
Cs 2O6
Li2O24
Na2O22
K2O21
Rb2O21
Cs 2O21
NaO23
KO22
RbO22
CsO22
NaO35
KO35
RbO35
CsO35
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 H2 O2 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 M2 O2 + M → 2 M2 O or
2 MNO3 + 10 M → 6 M2O + N2)
Properties:
oxides are stable up to 500 °C,
peroxides are stable up to 500 - 600 °C (Li 2O2 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 OHhyperoxides: 2 O2- + 2 H2O → O2 + H 2O2 + 2 OHozonides:
4 O3- + 2 H2O → 5 O2 + 4 OHApplication: Na2O2 – leaching agent
Li2O2, Na2O2, KO2 – oxygen source and CO2 absorber in respirators
(e.g. 4 KO2 + 2 CO2 → K2 CO3 + 3 O2)
General Properties of Alkali Metals (4)
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:
application:
electrolysis of molten LiCl/KCl mixture at 400-460 °C
- synthesis of Li hydride and Li amide
- synthesis of organic lithium compounds
(reducing age nt, polymerisation catalysts)
- manufacture of extremely light
and strong Al -Li alloys (2 - 3 % Li)
for space applications
- batteries
Lithium carbonate (Li 2CO3)
-
manufacturing:
-
application:
precipitation from soluble Li salts
2 Li + + CO32- → Li2CO3
- 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:
application:
Li2CO3 + Ca(OH) 2 → CaCO3 + 2 LiOH
- manufacturing of Li soaps and greases
Lithium hydride (LiH)
-
manufacturing:
application:
2 Li + H2 → 2 LiH at 700 °C
- drying agent
- hydrogen storage
- reducing agent in organic chemistry (LiAlH4, LiBH3)
Lithium nitrate (LiNO3)
-
manufacturing:
-
application:
Li2CO3 + 2 HNO3 → 2 LiNO3 + CO2 + H2O
in aqueous solution
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+ + HCO3NH4+ + 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
(CaCl 2 cannot used in further processes and is an waste
difficult to depose)
Summary:
2 NaCl + CaCO3 → Na2CO3 + CaCl 2
(occurs in aqueous solution in the opposite direction)
Sodium and its Salts
Salt
Manufacture
elementary
Na
- electrolysis of molten NaCl (modified
Downs process)
NaCl
Na2 CO3
(soda ash)
NaHCO3
NaNO3
Na2 SO4
(Glauber’s
salt)
NaHSO4
NaB4 O7
(borax)
Application
- reducing agent, catalyst in organic
chemistry
- manufacture of NaH, NaBH4 , Na2 O2
etc.
- coolant in nuclear reactors
(fast breeders)- sodium-sulphur
batteries
- mining or underground solving of
- starting material for all other inorganic
rock salt deposits
Na compounds
(in Germany - Staßfurt/Zielitz, Austria,
(e.g. Na2 CO3 , NaOH, Na2 SO4 ,
Spain, USA, Russia) and purification
Na2 B4 O7 , Na2 SiO 3 )
by flotation,
- raw material for chlorine alkali
- vaporising, freezing or electrodialysis
electrolysis
of sea water
- food industry
- mining of trona deposits (USA),
- glass industry
purification by solving, evaporating and
- synthesis of inorganic Na salts
calcination
- pulp and paper industry
- Solvay process (Europe)
- soap and detergent production
2 NaCl + CaCO3 → Na2 CO3 + CaCl2
- food industry (baking powder
- Na2 CO3 +H2O + CO2 → 2 NaHCO3
production)
(high purity)
- animal feedstuff
- rubber, chemical, pharmaceutical,
textile, leather and paper industries
- mining of natural deposits (Chile)
- fertilizer
- Na2 CO3 +2 HNO3 → 2 NaNO3 + H2O
+ CO2
- mining of natural deposits (Russia,
- pulp and paper industry
USA, Canada)
- additive in detergents
- glass industry
- 2 NaCl + H2 SO4 → Na2 SO4 + HCl
chemical
industry
(800 °C)
- 2 NaCl + MgSO4 → Na2 SO4 + MgCl2
(in aq. solutions)
(deep temperature precipitation of
Na2 SO4 )
- cleaning agents
- 2 NaSO 4 + H2 SO4 → 2 NaHSO4
- flux
- byproduct of CrO 3 manufacture
(Na2 Cr2O7 + 2 H2SO4 → 2 CrO3 + 2
NaHSO4 + H2 O)
- extraction from borate minerals
- glass, enamel, china and ceramic
(dissolving in H2 O, followed by
industries
selective crystallisation at 60 °C)
- manufacture of perborates for
- dehydratisation by calcination at 350detergents
400 °C
- flux, falme and corrosion inhibitor
Potassium and its Salts
Salt
Manufacture
elementary K
- KCl + Na → K + NaCl (760-880 °C,
favoured process)
- 2 KF + CaC 2 → CaF2 + 2 C + 2 K
(1000-1100 °C)
KCl
KBr
KI
K2 CO3
(potash)
KNO3
(saltpetre)
K2 SO4
Application
- only limited importance
- manufacture of KO2
- manufacture of low melting Na-K
alloys (reducing and drying agent, heat
transfer medium)
- mining or underground solving of salt
- starting material for all other
deposits (deposits in Germany inorganic K compounds
Staßfurt/Zielitz, Hanover, Werra/Fulda - production of K containing fertilizers
region - France, Canada, USA, Russia)
(KCl, K2 SO4 , KNO 3 )
and purification by flotation
- metallurgy, enamel industry,
manufacture of soaps
- manufacture of special IR transparent
optical glassware
- halogenation of potash with Fe(II, III)- photographic industry
bromide
- manufacture of special IR transparent
optical glassware
4 K2 CO3 + Fe3 Br8 → 8 KBr + Fe3 O4 +
4 CO2
- bromation of potash
3 K2 CO3 + 3 Br2 → 5 KBr + KBrO 3 + 3
CO2
- halogenation of potash with Fe(II, III)- photographic industry
iodide
- manufacture of special IR transparent
optical glassware
4 K2 CO3 + Fe3 I8 → 8 KI + Fe3O4 + 4
CO2
- reduction of KIO 3
- carbonation of KOH
- glass and enamel industry, pigment
manufacture
KOH + CO2 → K2 CO3 + H2 O
(precipitation, calcination at 250-350 - manufacturing of soaps and detergents
- food industry
°C)
- starting material for many inorganic
and organic K compounds
- fertilizer
- KCl + NaNO3 → NaCl + KNO3
- component of gun powder
- 2 KCl + 2 HNO3 + 1/2 O2 → 2 KNO3
+ Cl2 + H2 O
- 2 KCl + H2 SO4 → K2 SO4 + HCl (700 - fertilizer (trading of K2 SO4 and K2 SO4
°C)
· MgSO4 )
- 2 step process in aq. solutions:
(1) 2 KCl + 2 MgSO4 → K2 SO4 ·
MgSO4↓ + MgCl2
(2) K2 SO4 · MgSO4 + 2 KCl → 2 K2 SO4
+ 2 MgCl2