LECTURE SCHEDULE
Date
Topic
1.
Mon 31.10.
Review of Elements, Periodic Properties & Periodic Table
2.
Wed 02.11.
Short Survey of the Chemistry of Main Group Elements
3.
Fri
04.11.
Redox Chemistry (Mustonen)
4.
Fri
11.11.
Transition Metals: General Aspects & Crystal Field Theory (Thomas)
5.
Wed 16.11.
Ag, Au, Pt, Pd & Catalysis (Karttunen)
6.
Fri
Zn + Ti, Zr, Hf & Atomic Layer Deposition (ALD)
7.
Mon 21.11.
Cr, Mo, W & Metal Complexes
8.
Wed 23.11.
Mn, Fe, Co, Ni, Cu & Magnetism and Superconductivity (Mustonen)
9.
Mon 28.11.
Resources of Elements & Rare/Critical Elements &
Element Substitutions
18.11.
10. Wed 30.11.
Lanthanoids + Actinoids & Luminescence
11. Fri
Inorganic materials chemistry research
02.12.
EXAM ??
QUESTIONS: Lecture 2
1. Draw semitopological diagrams for the following boranes: B4H10 and B5H11
2. Which of the main groups is chemically most/least coherent?
MAIN GROUP ELEMENTS
n
Hydrogen: position in Periodic Table, isotopes, hydrides,
hydrogen storage
n
Alkali metals: group trends, Li-ion battery
n
Alkaline earth metals versus alkali metals
n
Boron group: crystal structures & melting points, boranes,
borides, BNCT
n
IsoelectroniC: C-C ja B-N
n
Nitrogen group: metal character and basicity of oxides
(N < P < As < Sb < Bi)
n
Multitude of sulphur compounds
n
Lightest element versus other group members: F – Cl, Br, I
n
Ionization energies and compounds of noble gases
Where would you place hydrogen in Periodic Table ?
1
18
2
13 14 15 16 17
3
4
5
6
7
8
9
10 11 12
HYDROGEN in the Periodic Table:
Group 1:
PRO:
- one s electron
- monovalent cation
CONTRA:
- not a metal
- does not react with water
Group 17: PRO:
CONTRA:
Group 14: PRO:
- one electron to the full shell
- non-metal
- two-atom molecule
- anion not common
- relatively inert
- half-filled electron shell
- electronegativity 2.2
Hydrogen, Deuterium & Tritium
n
H2 99.985%, D2 0.015 %, T2 10-15 %
n
Tritium is radioactive (half-life-time 12 years) but found in atmosphere upon cosmic radiation
n
Industrial preparation: in nuclear reactors from 6Li; bound into metals, e.g. UT3
n
Uses: medical application as a radioactive tracer, hydrogen bombs, fusion reactors
n
H, D & T: all physical properties different, e.g. boiling point
n
H, D & T: also chemical properties slightly different, e.g. covalent bond strength
n
H-O bond weaker than the D-O bond (gets broken more easily)
® electrolysis of water is utilized for the separation of D2O from H2O
HYDRIDES
n
Binary compounds of hydrogen
n
Hydrogen forms hydrides with most of the elements
n
Electronegativity of hydrogen only little higher than the average
electronegativity of elements → many of the ”hydrides” do not
contain the H- hydride ion
n
Hydrides are categorized according to the type of bonding:
ionic, covalent or metallic
n
IONIC HYDRIDES
- with alkali and alkaline earth metals (except Be, Mg)
- metal cation and H- ion
- crystal structures similar to those of halides (Cl replaced by H)
- very reactive, e.g.: LiH + H2O ® LiOH + H2
(used as reductants)
n
COVALENT HYDRIDES
- with all nonmetals (except noble gases) and the most electronegative metals (Sn,Ga)
- most of these hydrides are gaseous at room temperature
METALLIC (TRANSITION METAL) HYDRIDES
§
§
§
§
§
§
§
§
§
solid materials
hydrogen in interstitial positions
often nontoichiometric, e.g. TiH1.9
volume increases upon hydrogen intercalation
electrical conductivity decreases
brittle
used for e.g. storage of hydrogen
Ti(s) + H2(g) + heating/pressure ® TiH1.9(s)
TiH1.9(s) + higher temp. ® Ti(s) + H2(g)
ALKALI METALS
n
Chemically highly coherent group, but lithium somewhat
more different (diagonal relationship Li-Mg)
n
Metal radius (Å):
Li 1.52 (Mg 1.60), Na 1.86, K 2.27, Rb 2.48, Cs 2.65
n
Ionic radius (Å):
Li 0.76 (Mg 0.72), Na 1.02, K 1.38, Rb 1.52, Cs 1.67
n
Charge/ionic radius:
Li 1.40, Na 0.88, K 0.66, Rb 0.60, Cs 0.55
n
Melting point (oC): Li 180, Na 98, K 64, Rb 39, Cs 29
n
Oxidation product: Li2O, Na2O/Na2O2, K2O2, KO2, RbO2, CsO2
n
Li compounds more covalent than others
n
Li compounds dissolve more easily into nonpolar solvents,
and less into water
n
Only Li forms the nitride, Li3N (ref. Mg3N2)
n
Li salts often contain water of crystallization,
e.g. LiClO4·3H2O (ref. MgClO4·6H2O)
LITHIUM-ION BATTERY
n
Lithium: the lightest of all metals
the greatest electrochemical potential
the largest energy density per weight
small and easy/fast to move
n
Charging: Li-ions from cathode to anode
Discharging: Li-ions from anode to cathode
n
Cathode: Lix(Co/Ni)O2 (Goodenough 1980)
n
Anode: graphite
n
Commercialization: Sony 1991
n
Used: portable electronics
T. Motohashi, et al.,
Chemistry of Materials 19, 5063 (2007).
SAFE CHOICES FOR ELECTRODE MATERIALS
Li1-xFePO4 CATHODE (pos. electrode)
n
n
n
n
Olivine structure (not layered)
Safe, cheap, environmentally benign
® excellent for large-scale applications
Stable, constant voltage (3.4 V),
excellent cyclability, fast to charge
Problems: low electrical conductivity,
water-related stability issues
Li4+yTi5O12 ANODE (neg. electrode)
§ Spinel structure Li[Li1/6Ti5/6]2O4
§ No protective SEI (solid-electrolyte
interface) layer formed/needed ® safe
§ Stable, constant voltage (1.55 V), no
volume change, fast to charge
§ Problems: low electrical conductivity,
water-related stability issues
traditional
cathode
LiFePO4
LiCoO
Li4Ti5O12
ALKALI VERSUS ALKALINE EARTH METALS
n
Alkaline earth metals have larger Zeff than alkali metals
® smaller
® denser
® harder
n
Alkaline earth metals have two valence electrons per atom
® metal bonds stronger
® better electrical conductivity
® higher melting and boiling points
Li/Be
Na/Mg
K/Ca
Rb/Sr
Cs/Ba
Electronegativity
1.0/1.5
0.9/1.2
0.8/1.0
0.8/1.0
0.7/0.9
Metal radius (Å)
1.52/1.12
1.86/1.60
2.27/1.97
2.48/2.15
2.65/2.22
Density (g/cm3)
0.53/1.85
0.97/1.74
0.86/1.55
1.53/2.63
1.87/3.59
Melting point (oC)
181/1289
98/650
64/842
40/769
28/729
RT-resistivity (mohm cm)
9.47/3.70
4.89/4.48
7.39/3.42
13.1/13.4
BORON GROUP: MELTING POINTS
2180 °C
660 °C
30 °C
157 °C
303 °C
Boron
Aluminum
Gallium
Indium
Thallium
Rhombohedral
CCP
Orthorhombic
Tetragonal
HCP
BORON: ELECTRON DEFICIENT COMPOUNDS
n
B: 1s2 2s2 2p1
3 valence electrons ® octet not possible with covalent bonds
electronegativity 2.0 ® ionic bonds not possible
n
Difficult to achieve electron octet ® “electron deficient”
n
Hydrogen compounds of boron, boranes, are typical examples of
electron deficient compounds
n
Also Al, Ga, Be and Li may form electron deficient compounds:
(AlH3)n, Al2(CH3)6, [Be(CH3)2]n, [Li(CH3)]4
n
Normal single bond: 2 atoms and 2 bond electrons (2c-2e)
n
Electron deficient compound: (3c-2e)
BORANES
n
BnHn+4 or BnHn+6 (n = 2 - 20)
n
First synthesized by Alfred Stock 1909:
MgB2 + HCl ® B4H10, B5H9, B5H11, B6H10, B10H14, …
n
Air-sensitive compounds
® closed system, inert synthesis atmosphere
n
Naming:
number of B atoms with prefix,
number of H atoms in parenthesis at the end,
e.g. B2H6 diborane(6), B4H10 tetraborane(10), B5H9 pentaborane(9)
n
(3c-2e) weak bond ® breaking the bond requires little energy
B-O bond strong ® formation releases lot of energy
® burning of boranes is strongly exothermic
(1950s boranes were investigated as “superfuels” in aircrafts and spaceships)
SEMITOPOLOGICAL DIAGRAMS
n
William Lipscomb (Nobel 1976)
n
Bonds:
B-H (terminal 2c-2e)
B-H-B (3c-2e)
B-B (2c-2e)
B-B-B (3c-2e)
H
EXAMPLE:
B5H9
Valence electrons:
5 x 3e + 9 e = 24 e
Bonds:
5
4
2
1
2c-2e
3c-2e
2c-2e
3c-2e
B-H
B-H-B
B-B
B-B-B
10 e
8e
4e
2e
BORIDES
n
Binary compounds of boron
- more than 200 different compounds
- stoichiometry varies M5B … MB66
(e.g. M2B, MB, MB2, MB4, MB6)
- M-rich typical for d-block transition metals,
B-rich for main group metals, lanthanoids and actinoids
- know are also nonstoichiometric and mixed borides
- extremely hard, high-temperature resistive, chemically inert
- uses: coatings, electrodes, nuclear technology (protection, neutron
counter: 1n + 10B ® 7Li + 4He)
n
Boron carbide
- “B4C” = B12C3
- one of the hardest materials
- uses: armor material, bycycles
BORIDE STRUCTURES
n
M-rich:
(1)
(2)
(3)
(4)
(5)
separate B atoms
separate B2 atom pairs
Bn chains
double chains
planar nets (MB2)
(catenation tendency of B atoms increases with increasing B content)
n
B-rich:
(1) planar nets
(2) 3D nets
(3) clusters (e.g. B6 octahedral and B12 icosahedra)
TiB2
MgB2
LaB6
BONDING in BORIDES
n
The simple ionic / covalent / metal bond concepts fail
n
LaB6: - in B6 clusters covalent (multicenter-2e) bonds
- between clusters covalent (2c-2e) B-B bonds
- move of two electrons from La to B6 cluster
® La2+-B62- ionic bond
- third valence electron of La is delocalized (metal bond)
® good electrical conductivity, good electron emission
® use: electron guns in electron microscopes
Resistivity vs. temperature
Akimitsu 2001:
Superconductivity in MgB2
with Tc at 39 K
Crystal structure of MgB2
Magnetic susceptibility
vs. temperature
BORON-NEUTRON-CAPTURE-THERAPY (BNCT)
n
Possible treatment in particular for inoperable brain tumors/cancer
n
Boron (10B) absorbs efficiently neutrons (large absorption cross section)
→ radioactive 11B:a
n
Radioactive radiation destroys tumor cells
n
Challenge: to deliver high concentrations of a boron compound
specifically to tumor cells
n
Some boron compounds such as (B12H11SH)2- ion based ones have a
tendency to accumulate in cancer cells
10B
Isotope
Neutron-abs.
cross-section
(barn)
1 barn = 10-24 cm2
+ 1n ®
11B
Hydrogen-1
Boron-10
0.33
3.8 x 103
® 7Li + 4He
Carbon-12 Nitrogen-14
3.4 x 10-3
1.8
Oxygen-16
1.8 x 10-4
ISOELECTRONIC: C-C and B-N
n
Catenation (= chain-formation capability) typical for the carbon group
(ref. organic chemistry)
n
Catenation capability decreases: C > Si > Ge > Sn > Pb
(after C the next in catenation capability is S)
n
Also in the elemental form the C-C bond is strong, ref. carbon
allotropes:
graphite, diamond, fullerenes, nanotubes and graphene
n
Boron has one valence electron less and nitrogen one valence
electron more than carbon
® Boron nitride BN has in average the same number of valence electrons
per atom as carbon
® C ja BN isoelectronic
® resemble each other
DIAMOND
n
Diamond and graphite known for thousands of years
n
The fact that they are different forms of the same element was understood
much later
n
The critical conbustion experiments were made in the end of 18th century:
the same mass of diamond and graphite produces the same amount of CO2
n
In nature diamonds have been formed slowly from graphite under highpressure high-temperature conditions
n
The first synthetic diamonds were made in 1953 in Sweden under 7 GPa
(= 70 000 atm) and at high temperature using molten Ni as a catalyst
(nowadays most of the diamonds are man-made)
n
In diamond the C atoms are packed much more denser (3.5 g/cm3) compared
to the graphite (2.3 g/cm3); this is the reason why the high-pressure conditions
promotes the formation of diamonds
n
For high-tech applications diamond thin films are needed; these are
synthesized using hydrocarbons as precursors (methanol, ethanol)
(Note: the same tetrahedral sp3 hybridization as in diamond)
HIGH-PRESSURE
SYNTHESIS
5 GPa = 5 x 104 atm
HIGH-PRESSURE SYNTHESIS
•
•
•
•
5 GPa = 50 000 atm
400 – 1200 oC
10 – 120 min
50 – 100 mg
HP equipment
at Tokyo Tech
H. Yamauchi & M. Karppinen, Supercond. Sci. Technol. 13, R33 (2000).
HP equipment at Aalto
§ 4 GPa & 25 GPa
BORON NITRIDE
n
Synthesis of boron nitride: B2O3 + NH3 at 1000 oC ® graphite-type BN
n
High-pressure high-temperature treatment ® diamond-type BN (so-called boratson)
n
Also fullerene-type BN molecules known
DIAMOND STRUCTURE: C and BN
n
Each C atom (or B and N atom) bonded with strong covalent bonds to
tetrahedrally surrounding neighbours
n
In the 3D atom lattice each C-C bond is equally long (1.54 Å) and strong
n
All four valence electrons of carbon are bound in the four bonds
® pure diamond is one of the best known electrical insulators
n
Strong covalent bonds ® diamond is the strongest material known
n
Diamond-type BN is the second strongest material (ca. 50 % of the
harness of diamond)
n
Strong bonds ® high melting point (4100 oC)
DIAMOND AS THERMAL CONDUCTOR & SEMICONDUCTOR
n
Diamond is the best thermal conductor among the known materials
(ca. 5-times better than the best metals)
n
In diamond thermal conduction happens via phonons (vibrations of C atoms)
[in metals heat (like electricity) transport occurs via moving valence electrons]
n
C atoms are light and the C-C bond is strong
® efficient heat conduction
(impurities and lattice imperfections depress thermal conductivity)
n
BN » “impure C”
® diamond-type BN is not as good thermal conductor as diamond
n
Pure diamond is electrical insulator (used as heatsink for semiconductor laser diodes)
n
B-doping of diamond → p-type semiconductor
n
Diamond-type BN:
- slight B excess ® p-type semiconductor
- slight N excess ® n-type semiconductor
GRAPHITE STRUCTURE: C and BN
n
C: the fourth valence electron loosely bound ®
good electrical conductivity ® black
n
BN: Polar bond between the layers ® no electrical
conductivity ® white
METALLIC CHARACTER and BASICITY OF OXIDE: N < P < As < Sb < Bi
n
N (g), P (s; mp. 44 oC), As (s), Sb (s), Bi (s)
(N-N bond weaker than C-C bond due to the repulsion between from the
extra electron pair ® no catenation but :N º N: gas)
n
Oxidation states:
n
Acid-base nature of oxide reflects the nonmetal/metal character of the
element
- N oxides (+I ... +V), P2O3 and P2O5 acidic
- As2O3, As2O5 and Sb2O3 amfoteric
- Bi2O3 basic
n
N, P nonmetals, As, Sb semimetals, Bi metal
n
Melting points (oC):
N -210, P 44, As subl. 615, Sb 631, Bi 271
(for metals m.p. decreases in a group from up to down, for nonmetals the
trend is opposite)
n
Note: even though Bi most metallic, resistivity increases As < Sb < Bi
N: -III ... +V
P: -III, +III, +IV, +V
As, Sb, Bi: +III, +V (Bi: inert-pair effect)
MULTITUDE OF SULPHUR COMPOUNDS
n
Several stable oxidation states
n
Allotropy of S: tens of allotropes (ref. O 2, Se 6, Te 3, Po 2)
n
Various S-chains and S-rings, where other atoms can bond
n
Most common: S8-ring ”crowne” (different polymorfs)
n
Polycations Sn2+ (n = 4 yellow, n = 8 blue, n = 19 red)
n
Polysulfides Sn2- (n = 2 - 6; strongly colourful)
n
Thio-compounds
n
Sulphur is also one of the basic elements in organic compounds
HALOGENS
n
Chemically very homogeneous group,
except fluorine which differs from the
rest of the group members in many ways
Element
Melting
Point (°C)
Boiling
Point (°C)
F2
-219
-188
Cl2
-101
-34
Br2
-7
+60
I2
+114
+185
Wölsendorf, Germany
Antozonite (Stink Spar)
Elemental fluorine is generated through minute uranium
inclusions in the mineral, which constantly emit ionizing
radiation and thus split the fluorite into calcium and
elemental fluorine. The fluorine remains in minute
inclusions, separated from the calcium by the nonreactive fluorite and thus retains its elemental form.
Fluorine subsequently reacts with atmospheric
oxygen and water vapor, producing ozone (whose
characteristic smell, originally gave the name for
the mineral.
2012: First proof of F2 on Earth by 19F-MAS-NMR
J. Schmedt auf der Günne, M. Mangstl & F. Kraus, Angew. Chem. Int. Ed. 51, 7847 (2012).
FLUORINE VERSUS OTHER HALOGENS
n
Atomic radius (Å): F 0.71, Cl 0.99, Br 1.14, I 1.33, At ???
n
Electronegativity: F 4.0, Cl 3.0, Br 2.8, I 2.5, At 2.2
n
Dissosiation energy (kJ/mol): F 155, Cl 240, Br 190, I 149, At 116
n
Oxidation states:
F:
Cl:
Br:
I:
-I
-I, +I, (+II), +III, (+IV), +V, (+VI), +VIII
-I, +I, +III, (+IV), +V, (+VI), +VIII
-I, +I, +III, +V, +VIII
Small F atoms
® Large electron-electron repulsion
® Small bond energy
® Highly reactive
Preparation of fluorine (from fuorite CaF2) Moissan 1886 (Nobel 1906):
- 2 HF ® H2 + F2
- problem: F2 and HF gases highly reactive and poisonous
- electrolysis in KF melt, Pt cell, Pt-Ir electrodes, -50 oC
n
Fluorine obeys always the octet rule
(only one covalent bond possible)
n
Large electronegativity
® strong oxidizer
® promotes high oxidation states for cations *
® compounds ionic
® strong hydrogen bonds
* Cations often have higher oxidation states on
oxides compared to fluorides,
e.g. OsO4 possible, but not OsF8
Reason: 8 F- ions around Os8+ cation not possible !
NOBLE GASES: BASIC PROPERTIES
n
Helium has the lowest known boiling point (-269 oC)
n
Ionization energies (kJ/mol):
He 2269, Ne 2079, Ar 1519, Kr 1349, Xe 1169, Rn 1036
COMPOUNDS
n
Bartlett 1962: O2 (g) + PtF6 (g) ® [O2+][PtF6-](s)
n
Ionization energy of O2 1180 kJ/mol » IE(Xe)
® “[Xe][PtF6] “ ® Ar, Kr, Xe form compounds with the most
electronegative elements
n
Oxidation states: +II, +IV, +VI, +VIII
n
Bond strengths (kJ/mol): Xe-F 130, Xe-O 84, Kr-F 50
n
KMXeNaO6 double perovskite !
[S.N. Britvin et al., Angew. Chem. 127, 1 (2015)]
SUPERFLUIDS
n
Viscosity zero
n
In continuous movement,
escape from the container
n
Classical mechanical laws do
not work, described by
quantum mechanics
n
Lambda point:
temperature where materia
transforms to a superfluid
n
Known supra liquids: 3He and
4He
n
Superfluid He conducts heat
extremely well
n
Phenomenon discovered by
Heike Kamerlingh Onnes