Download Lecture 2: Bonding in solids

Lecture 8: Transition metals
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Oxidation states in transition metal
compounds
Some general trends and classifications
– Compounds with main group elements
Intermetallic compounds
Ligand field theory
– Fundamentals of magnetism in
transition metal compounds
(full treatment of magnetism on
lecture 9, metal oxides)
Case study: the missing Cu(I)F
Figure: Wikipedia
Figure: Müller
Figure: AJK
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Literature
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Transition metals (d-block)
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Many oxidation states, very rich chemistry
Many magnetic compounds (unpaired d-electrons)
Many colorful compounds due to d-d transitions
Figure: Wikipedia
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Elemental transition metals are
close-packed
Cu (fcc)
Legend:
.. / .. = mixed structure
[...] = predicted
structure
P
I (bcc)
F (fcc)
Fe (bcc)
Figures: AJK
Zn(hcp)
Figure: Wikipedia
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Electron configurations (3d metals)
Ref: chemwiki.ucdavis.edu
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Oxidation state
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The oxidation state of a transition metal is a key concept for understanding the
chemistry of transition metal compounds
Indicates the degree of oxidation (loss of electrons) of an atom in a compound
For example Fe:
– FeCl2 -> Cl– and Fe2+ / Fe(II) -> iron(II) chloride
– FeCl3 -> Cl– and Fe3+ / Fe(III) -> iron(III) chloride
The exact definition of an oxidation state is actually still debated
IUPAC Technical Report Toward a comprehensive definition of oxidation state:
– P. Karen et al. Pure Appl. Chem. 2014, 86, 1017–1081
– “The oxidation state of a bonded atom equals its charge after ionic
approximation”
– “In the ionic approximation, the atom that contributes more to the bonding
molecular orbital (MO) becomes negative”
The sum of oxidation states in a compound must be zero!
The oxidation state of an atom in a compound is a useful concept, but it is really
just a concept. In particular, the “charges” are not real!
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Known oxidation states for d-block
Transition metals show a larger
number of oxidation states in
comparison to main group elements
Ref: https://en.wikipedia.org/wiki/
List_of_oxidation_states_of_the_elements
(based on Greenwood and Earnshaw +
recent literature)
Year
2014!
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Nature 2014, 514, 475.
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d-block + p-block compounds (1)
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d-block + group 17 (F, Cl, Br, I)
– All transition metals form halides with group 17 elements
– Many are molecular, but numerous polymeric / 3D structures are also known
– For example, ICSD lists 1025 binary transition metal halides (April 2016)
– Few of the 198 cubic binary transition metal halides are listed below
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d-block + p-block compounds (2)
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d-block + group 16 (O, S, Se, Te)
– Transition metal oxides show extremely rich chemistry, discussed on Lecture 9
– 1936 binary oxides, 13300 ternary oxides (April 2016)
– 1110 binary sulfides, 720 binary selenides, 642 binary tellurides
d-block + group 15 (N, P, As, Sb, Bi)
– Many nitrides are particularly important due to their hardness (coatings)
– E.g. TiN (rocksalt structure)
– Numerous phosphides, arsenides, and antimonides are known. Bismuth is
already so metallic that it rather forms alloys
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d-block + p-block compounds (3)
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d-block + group 14 (C, Si, Ge, Sn, Pb)
– ICSD lists 636 binary carbides (April 2016), e.g. Fe3C, cementite, present in
most steels.
– Some carbides such as tungsten carbide (WC) are extremely hard
– Organometallic chemistry (compounds with M-C bonds) is a huge branch of
chemistry, but since it mainly deals with molecular species, not discussed here
– Obvious exception: Metal-organic frameworks with M-C bonds
– Many silicides and germanides are well-known, but tin is so metallic that it
rather forms alloys (e.g. Cu + Sn = bronze)
d-block + group 13 (B, Al, Ga, In, Tl)
– ICSD lists 524 binary borides (April 2016)
– Similar to some carbides and nitrides, many borides are very hard materials
e.g. OsB2, WB4, and ReB2
– Al, Ga, In, Tl -> metallic, rather form alloys and intermetallic compounds
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The literature listed above
contains very detailed
accounts on various
compounds of transition
metals.
Naturally, new compounds
are found all the time
ICSD is nowadays also a
great resource for starting a
literature search
Greenwood-Earnshaw
Cotton and
Wilkinson
Holleman-Wiberg
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Intermetallic compounds
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Solid-state compound exhibiting
– Metallic bonding (metal-metal bonds)
– Defined stoichiometry
– Ordered crystal structure (different from metal alloys, where the atoms are
completely mixed)
Definition normally includes also s- and p-block metals (Na, Mg, Al, etc.)
– For example, Cu3Al, NiAl, Ni3Al, Cu3Sn… (ordered structures!)
Often offer a compromise between ceramic and metallic properties
Can also display desirable magnetic, superconducting, and chemical properties,
due to mixed metallic and covalent/ionic bonding
Examples discussed here:
– Laves phases
– Heusler compounds
Ref: Wikipedia
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Laves phases
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Laves phases form the largest group of intermetallic compounds
– Often also referred to as Friauf-Laves phases
General composition AB2 with the larger A atoms in the center of a 16-atom
(Frank–Kasper) polyhedron and smaller B atoms in the centers of icosahedra (12-c)
More than 1400 binary and ternary Laves phases are reported in “Pearson’s
Handbook of Crystallographic Data for Intermetallic Phases”
Three different polytypes: MgCu2 (cubic), MgZn2 (hex), and MgNi2 (hex).
Mg sublattice
Ref: Stein, Intermetallics 2004, 12, 713–720
MgCu2 (Fd-3m)
Related to diamond structure
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Heusler compounds
Cu
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Heusler phases are intermetallics with
particular composition and face-centered
cubic crystal structure
– Full Heusler: X2YZ (e.g. Cu2MnAl)
– Half-Heusler: XYZ (e.g. NiTiSn, CoTiSb)
Mn
Al
Figures: AJK
Cu2MnAl (Fm-3m)
Metallic, ferromagnetic
Ti (8-coord)
Sn
Ni
Figure: Wikipedia
NiTiSn (F-43m)
Semiconductor (18 VE!), non-magnetic
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Ligand field theory (1)
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The mutual interaction between bonding electron pairs is the same for transition
metal compounds as for compounds of main group elements
However, nonbonding valence electrons behave differently
For transition metal atoms these generally are d electrons that can be
accommodated in five d orbitals.
In what manner the electrons are distributed among these orbitals can be judged
with the aid of ligand field theory (LFT)
LFT considers how the d electrons have to be distributed so that they attain a
minimum repulsion with each other and with the bonding electron pairs
In its original version by Hans Bethe, it was formulated as crystal field theory (CFT)
– Consider only electrostatic repulsion between the d electrons and the ligands,
which are treated as point-like ions
The ligand field theory is in principle superseded by the molecular orbital theory,
but it still is a convenient and simple tool to understand the bonding and
magnetism of transition metal compounds
Ref: Müller p. 73
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CFT
Degeneracy breaks
Electrostatic repulsion
between the
surrounding anions
and metal d-electrons
Figures: Chemwiki
d-electrons still
degenerate
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Ligand field theory (2)
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In ligand field theory, the ligands are not just considered as negative point charges,
but the covalent nature of the sigma-bonding is also taken into account
The basic concept of the d-orbital splitting remains
Ref: Wikipedia
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Ligand field theory (3)
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Ref: Müller p. 74
In an octahedral geometry, The energy difference between the occupation of a t2g
and an eg orbital is termed ΔO
The value of ΔO depends on the repulsion exercised by the bonding electron pairs
on the d electrons
Compared to a transition metal atom the bonded ligand atoms are usually much
more electronegative
– The centers of charge of the bonding electron pairs are much closer to them,
especially when they are strongly electronegative.
– Therefore, one can expect a decreasing influence on the d electrons and thus
a decrease of ΔO with increasing ligand electronegativity
– Decreasing ΔO values also result with increasing sizes of the ligand atoms (the
electron pairs are distributed over a larger space -> less repulsion with d-el)
In the presence of multiple bonds between the metal atom and the ligands, as for
example in metal carbonyls, the electron density of the bonds is especially high
and their action is correspondingly large
ΔO can be measured with spectroscopic methods and the spectrochemical series is
obtained by ordering different ligands according to decreasing ΔO :
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Ligand field theory (4)
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Ref: Müller p. 75
LFT can be used to explain the magnetism of transition metal compounds
Because of the splitting of the d-orbitals, there is an energetical competition
between high-spin and low-spin configuration of the d-orbitals
The first three d-electrons occupy the t2g orbitals in accordance to Hund’s rule
When four nonbonding electrons are present, there are two alternatives for the
placement of the fourth electron
– The energy necessary to include a second electron in an already occupied
orbital is called the electron pairing energy P
– If P > ΔO, the fourth electron will occupy an eg orbital (high-spin)
– If P < ΔO, the fourth electron will occupy a t2g orbital (low-spin)
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Jahn-Teller effect
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Ref: Müller p. 76
In a high-spin d4 complex only one of the two eg orbitals is occupied
If it is the dz2 orbital, then it exerts a strong repulsion on the bonding electrons of
the two ligands on the z axis
These ligands are forced outwards and the coordination octahedron suffers an
elongation along the z axis. This effect is known as the Jahn–Teller effect
Instead of the dz2 orbital the dx2-y2 orbital could have been occupied which would
have produced elongations along the x and y axes
– However, more energy would be needed to stretch four bonds
The Jahn-Teller effect is always to be expected when degenerate orbitals are
unevenly occupied with electrons. In fact, it is observed for the following
electronic configurations (in principle also for d1, but not in practice)
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Other coordination geometries
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The energetics and the splitting of the d-orbitals are different for other
coordination geometries. The figure below shows the most typical geometries for
transition metals
A group theoretical resource: http://symmetry.jacobs-university.de/
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Most common
coordination
polyhedra for CN 2-6
Ref: Müller p. 81
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Experimental-theoretical study on
the missing copper(I) fluoride
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The missing copper(I) fluoride
CuCl (F-43m)
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Existing copper halides:
Cu(I)
Cu(II)
CuF2
CuCl
CuBr
CuCl2
CuBr2
CuI
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CuCl, CuBr and CuI: Zinc blende (sphalerite)
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Aqueous system:
2 Cu+
→ Cu2+ + Cu
Liquid NH3 system:
Cu2+ + Cu → 2 Cu+
Approach of Florian Kraus:
– CuF2 + Cu → 2 CuF (in liquid NH3)
CuF2 (P 1 21/n 1)
Jahn-Teller
distorted Cu(II)
Cu-F (eq): 1.9 Å
Cu-F (ax): 2.3 Å
Figures: AJK
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Running the Cu2+ + Cu reaction in
liquid ammonia (–40°C)
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2 Cu + 2 CuF2 + 14 NH3 → “(CuF)4 ∙ 14 NH3” (colorless crystals)
{[Cu(NH3)3]2[Cu2(NH3)4]}F4 ∙ 4 NH3
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Ligand-unsupported cuprophilic
Cu…Cu interaction
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{[Cu(NH3)3]2[Cu2(NH3)4]}F4 ∙ 4 NH3
[Cu2(NH3)4]F2 chain
Metallophilic interactions: Weak metal-metal bonding
Best known example: aurophilicity (Au-Au, comparable to H-bond)
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Warming up “(CuF)4 ∙ 14 NH3” from
–40°C to room temperature
(CuF)4 ∙ 14 NH3 → 4 [Cu(NH3)2]F ∙ NH3 + 2 NH3
“CuF ∙ 3.5 NH3”
“CuF ∙ 3 NH3”
[Cu(NH3)2]F ∙ NH3
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Remove ammonia from “CuF ∙ 3 NH3”?
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22 hypothetical bulk CuF phases were investigated using quantum
chemical methods (by starting from various binary phases in ICSD)
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Previous studies on CuF phases
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F–Cu–F chains packed in the chiral cinnabar structure (P3121)
Cinnabar = mercury(II) sulphide, HgS
Cinnabar (HgS)
Figure: Wikipedia
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19.4.2016
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Antti Karttunen
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Relative energies of the CuF phases
Relative energies of CuF phases in kJ/mol per CuF unit
DFT-PBE0/TZVP level of theory
Structure
Space
group
Relative
energy
Cinnabar
P3121 (152)
0.0
Distorted wurtzite
Cmcm (63)
1.2
Wurtzite
P63mc (186)
3.9
Rocksalt
Fm-3m (225)
7.6
Sphalerite
F-43m (216)
10.7
Distorted structure
from the lowenergy vibrational
mode
Low-energy
vibrational mode
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Cinnabar vs. Distorted Wurtzite (Cmcm)
Chiral
Not
chiral
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Low-temperature route to
copper nitride Cu3N
Cu3N (Pm-3m)
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Normally, Cu3N is prepared in a temperature range of 250–325°C
from CuF2 in a stream of NH3
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The first evidence of chemical nitride formation at room
temperature
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Summary of copper(I) fluoride
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Binary CuF was not yet synthetized, but several new ammine Cu(I)-fluorides were
prepared and characterized
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One compound, {[Cu(NH3)3]2[Cu2(NH3)4]}F4 ∙ 4 NH3, shows rare ligandunsupported cuprophilic interaction
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In addition to CuF-cinnabar with helical F–Cu–F chains, an orthorhombic phase
with linear F–Cu–F chains is energetically feasible
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The experimental work continues (Prof. Florian Kraus, University of Marburg)
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