1) Anorganische C Metalls

2) The electron as a wave - the Schrödinger equationi) H’w = Ew

(a) H’: Hamilton operator(b) E: energy eigenvalue(c) w: wave function

ii) Probability of presence of the electron: -> elektrondensity ~ w2

3) Quantum Numbersi) Each of the wavefunctions obtained by solving the Schrödinger equation for a

hydrogeni atom is uniquely labeled by a set of three integers called the quantum numbers

(a) n Principal quantum number(b) l orbital angular momentum quantum number(c) ml magnetic quantum number(d) ms spin quantum number

4) Energy level scheme of the hydrogen atom

5) 3-Dimensional orientation of the orbitals

6) Electron spin

a) Two electrons are paired if they have directions with opposite spins (clockwise, counterclockwise)

b) Two electrons have parallel spins if their spins are oriented in the same direction

7) Quantum State

8) Energy level scheme of a multiple electron system

i) Used to predict the ground-state electron configurations by accommodating electrons in the array of molecular orbitals and recognizing the constraints of the Pauli principle.

ii) Filling rules (building-up principles):1. Pauli Principle: no electrons with the same set of quantum

numbers2. Energy sequence/order of the orbitals

3. Hund’s Rule: When more than one orbital has the same energy, electrons occupy separate orbitals and do so with parallel spins (↑↑).

4. Each orbital can accommodate up to two electrons. if more than one orbital is available for occupation (because they happen to have identical energies, as in the case of pairs of π orbitals), then the orbitals are occupied separately. In that case, the electrons in the half-filled orbitals adopt parallel spins (↑↑), just as is required by Hund’s rule for atoms

iii) With very few exceptions, these rules lead to the actual ground-state configuration of the Period 2 diatomic molecules. For example, the electron

configuration of N2, with 10 valence electrons, is N2:

9) Pauli Principle, Hund’s Rulei) Pauli Principle

(a) All the electrons of an atom must differ at least in one quantum numberii) Hund’s Rule

(a) All degenerate orbitals (thus all those with the same energy) must initially simply be occupied with electrons with same direction spins

10)Merkschema for the occupancy of the orbitalsi) Aufbau Principle (also the building-up principle)

11)The Electrochemical Series

i) A reaction is thermodynamically favorable (spontaneous) in the sense K>1, if E° >0, where E° is the difference of the standard potentials corresponding to the half-reactions into which the overall reaction may be divided.

ii) The oxidized member of a couple is a strong oxidizing agent if E° is positive and large; the reduced member is a strong reducing agent if E° is negative and large

iii) List is arranged: Ox/Red couple with strongly negative E° [Red is strongly reducing] … Ox/Red couple with strongly positive E° [Ox is strongly oxidizing]

12)The differentiation of metals in the Periodic Table of Elements

i) The classification of the elements: The elements are broadly divided into metals, nonmetals, and metalloids according to their physical and chemical properties

(a) Metals are typically lustrous, malleable, ductile, electrically conducting solids

(b) Nonmetals are often gases, liquids, or solids that do not conduct electricity appreciably

(c) Metalloids are elements with properties making it difficult to classify as metals or nonmetals

13)Electron configuration of a few metals

14)Structures of Metalsi) The close packing of spheres: close packing of identical spheres can result in a

variety of polytypes, of which hexagonal and cubic close-packed structures are the most common

(a) Many metallic and ionic solids can be regarded as constructions of hard spheres, free to pack together as closely as geometry allows, if no directional covalent bonding, adopting a close-packed structure (least unfilled space)

1. Coordination number (CN) of sphere in close-packed arrangement (‘number of nearest neighbors’) is 12

i. When directional bonding important, no longer close-packed and CN less than 12

ii) Densest packing (fcc & hcp): cubic face-centered lattice(a) Any collection of identical atoms, such as those in the simple picture of an

elemental metal, or of approximately spherical molecules, likely to adopt close-packed structure

1. Unless energetic reasons (covalent interactions) intervene

(b) face-centered cubic (fcc or ccp) spheres of the third layer are placed above the gaps in the first layer. The second layer covers half the holes in the first layer and the third layer lies above the remaining holes. this arrangement results in an ABCABC… pattern, where C denotes a layer that has spheres not directly above spheres of the A or the B layer

positions (but they will be directly above another C type layer). This pattern corresponds to a structure with a cubic unit cell

(c) hexagonally close-packed (hcp) spheres of the third layer lie directly above the spheres of the first. This ABAB… pattern of layers, where A denotes layers that have spheres directly above each other and likewise for B, gives a structure with a hexagonal unit cell

iii) body-centered cubic (bcc)

15)Overview of the structures of metals (at room temperature)

16)Densest Sphere-Packing

17)Hexagonal- and Cubic-Densest Packing

18)Elementary cells and sequence of layersi) hexagonally close-packed (hcp)

ii) face-centered cubic (fcc)

19)Hexagonally close-packeda) hexagonal elemental zellb) coordination number (CN), each atom has 12 neighbors at the same distance

20)Face-centered cubic packing

21)Face-centered cubic packinga) face-centered cubic elemental cellb) CN 12, the layers of the densest packing lie vertically to the body/space diagonal of

the elemental cell

22)Cubic Body-centered latticei) body-centered cubic (bcc or cubic-I): one commonly adopted arrangement has

the translational symmetry of the body-centered cubic lattice, in which a sphere is at the center of a cube with spheres at each corner

(a) metals with this have CN of 8ii) all alkali metals, Group 5, and Group 6 metals under standard conditions adopt

bcc structure

23)Occupancy of the energy bands of Li

24)Occupancy of the energy bands of Be

25)Band Model analogy

26)Classification according to position in the periodic table

27)Ionic crystal: gaps in fcc

a) Octahedron gaps

b) Tetrahedron gaps

28)Ionic crystal: AB structuresa) NaCl

b) CsCl

c) ZnS; d) ZnS

29)Sodium Chloride (NaCl) crystal

30)Lithium Chloride crystal

31)Cesium Chloride crystal

32)Zinc-blende (sphalerite) crystala) The sphalerite (zinc-blende) structure

b) The sphalerite (zinc-blende) structure projection representation. Note its relation to the ccp lattice with half the tetrahedral holes occupied by Zn2+

ions.

33)Zinc-blende: fcc S2-

34)Wurtzite crystali) The wurtzite structure takes its name from another polymorph of zinc sulfideii) Wurtzite differs from sphalerite structure in being derived from an expanded

hcp anion array rather than a ccp array, but as in sphalerite the cations occupy half the tetrahedral holes; that is just one of the two types (either T or T’)

(a) This structure, which has (4,4)-coordination, is adopted by ZnO, AgI, and one polymorph of SiC, as well as several other compounds

iii) local symmetries of cations and anions are identical with respect to their nearest neighbors in wurtziteand sphalerite but differ at the second-nearest neighbors

35)Wurtzite: hcp S2-

36)Ionic crystals: AB2 Structuresa) CaF2 (fluorite)

b) TiO2 (rutile)

c) SiO2

37)Fluorite Crystali) In fluorite, the Ca2+ ions lie in an expanded ccp array and the F- ions occupy all

the tetrahedral holesii) The cations are close-packed, because the F- anions are smalliii) lattice has (8,4)-coordination, consistent with there being twice as many anions

as cations(a) Anions, in their tetrahedral holes have four nearest neighbors(b) Cation site surrounded by a cubic array of eight anions

iv) Note relation of this structure to the caesium-chloride structure, in which all the cubic holes are occupied

38)Rutiles Crystali) The rutile structure (TiO2) takes its name from rutile, a mineral form of

titanium(IV) oxideii) Structure can be considered an example of hole filling in an hcp anion

arrangement, but now cations occupy only half the octahedral holes and there is considerable buckling of the close-packed anion layers

(a) Arrangement results in a structure reflecting strong tendency of a Ti4+ ion to acquire octahedral coordination

iii) Each Ti4+ atom is surrounded by six O atoms and each O atom is surrounded by three Ti4+ ions

(a) hence, rutile structure has (6,3)-coordinationiv) Principal ore of tin, cassiterite (SnO2), has the rutile structure, as do many metal

difluorides

39)The Group 1 elementsa) All group 1 elements are metallic but, unlike most metals, they have low densities

and are very reactive(1) Lithium (Lithium, Li)(2) Sodium (Natrium, Na)(3) Potassium(Kalium, K)(4) Rubidium (Rubidium, Rb)(5) Caesium (Caesium, Cs)

b) Form simple ionic compounds, most of which are soluble in water

c) Trends can be explained in terms of variations in atomic radii and ionization energies

40)Potassium chloride (Sylvin, KCl)i) Occurs naturally as the mineral sylvite and in combination with sodium chloride

as sylvinite

41)Occurrence and Extraction of Group 1 elements: Down’s process (Na production)i) The group 1 elements can be extracted by electrolysis

(a) Natural abundance of lithium is low, most abundant minerals being spodumene (LiAlSi2O6) and lepidolite, from which Li is most commonly extracted

ii) Sodium occurs as mineral rock salt (NaCl) and in salt lakes and seawateriii) NaCl deposits may be mined in the conventional way or water may be pumped

underground to dissolve rock salt, which is then pumped out as saturated brine solution

(a) Metal extracted by Down’s process, the electrolysis of molten sodium chloride:

iv) 2 NaCl(l) 2 Na(l) + Cl2(g)

1. Sodium chloride kept molten at 600°C, below melting (808°C), by addition of calcium chloride (CaCl2)

2. High potential difference (4-8 V) applied between a carbon anode and an iron cathode immersed in the molten salt

3. Electrolysis liberates liquid sodium metal at the cathode, which rises to the surface of the cell, where collected under inert atmosphere

4. Also used for industrial production of chlorine, which is generated at the anode

42)Lithium-ion Batterya) Very negative standard potential and low molar mass of lithium make it an ideal

anode material for batteriesi) have relatively high specific energy (energy production divided by the mass of

the battery) because lithium metal and compounds containing lithium are light in comparison with some other materials used in batteries (eg lead and zinc)

b) Lithium rechargeable battery mainly uses Li1-xCoO2 (x < 1) as the cathode with a lithium/graphite anode, LiC6.i) Lithium ions are produced at the anode during the battery dischargeii) To maintain charge balance, Co(IV) is reduced to Co(III) in the form of LiCoO2 at

the cathode(1) The reactions occurring during battery discharge are:

Cathode Li1-xCoO2(s) + xLi+(sol) + xe- LiCoO2(s)Anode C6 6 C(graphite) + Li+(sol) + e-

iii) Battery is rechargeable because both the cathode and the anode can act as host for the Li+ ions, which can move back and forth between them when charging and discharging

LiCoO2(s) + C(s) Li@C(s) + CoO2(s)iv) Solid-state: Li@C(s) indicates that a Li+ ion has penetrated between the

graphene sheets of graphite to form and intercalation compoundv) The reaction takes place in a lithium-ion battery during charging and its reverse

takes place during discharge.

43)18-Crown-6 ([K(18-Krone(6)]+, (C2H4O)6)i) Crown ethers form complexes with alkali metal ions that are reasonably stable in

nonaqueous solutions

44)Chloralkali processa) Chlorine and dihydrogen are the gaseous products

b) Industrial process is based on the electrolysis of aqueous sodium chloridei) Water is reduced to hydrogen gas and hydroxide ions at the cathode and

chloride ions are oxidized to chlorine gas at the anode:2 H2O(l) + 2e- H2(g) + 2 OH-

2 Cl-(aq) Cl2(g) + 2 e-

ii) There are three different types of cells that are used for the electrolysis:(1) In a diaphragm cell there is a diaphragm which prevents the OH- ions

produced at the cathode from coming into contact with the Cl2 gas produced at the anode(a) Made of polytetrafluoroethylene mesh(b) During electrolysis, solution at the cathode is removed continuously and

evaporated in order to crystallize the sodium chloride impurities(i) final sodium hydroxide solution contains ~1% by mass NaCl

(2) Membrane cell

(3) The mercury cell uses liquid mercury as the cathode(a) Chlorine gas is produced at the anode, but sodium metal is produced at

the cathode:Na+(aq) + e- Na(Hg)

(b) The sodium-mercury amalgam is reacted with water on a graphite surface:

2 Na(Hg) + 2 H2O(l) 2 NaOH(aq) + H2(g)(c) The sodium hydroxide solution produced by this route is very pure and

the mercury cell is the preferred source of high-quality, solid sodium hydroxide.(i) Unfortunately, the process accompanied by discharge of muercury

into environment

45) Sodium chloride structurea) The rock-salt structure is based on a ccp array of bulky anions with cations in all

octahedral holesi) Alternatively, can be viewed as a structure in which the anions occupy all the

octahedral holes in a ccp array of cationsb) stoichiometry AX: # of octahedral holes in a close-packed array equal to number of

ions forming the array (the X ions) and filling them with A ionsi) each ion surrounded by octahedron of six counter-ions, thus coordination

number of each type of ion is 6 and the structure is said to have(6,6)-coordination

c) The rock-salt structure

d) Its projection representationi) Note the relation of this structure to the fcc structure, with an atom in each

octahedral hole

46) Caesium chloride structurea) Much less common than the rock-salt structure for compounds of stoichiometry AX

is the caesium-chloride structurei) Possessed by CsCl, CsBr, and CsI, as well as some other compounds formed of

ions of similar radii to these, including TlI(1) Note, NH4Cl also forms this structure, despite relatively small size of NH4

+ ion because the cation can form hydrogen bonds with four of the Cl- ions at the corners of the cube

(2) Many 1:1 alloys, such as AlFe and CuZn have a caesium-chloride arrangement of the two metal atom types

b) Has a cubic unit cell with each corner occupied by an anion and a cation occupying the ‘cubic hole’ at the cell center (or vice versa)i) As a result, Z = 1

c) Alternative view of this structure is as two interlocking primitive cubic cells, one of Cs+ and the other of Cl-

d) The coordination number of both types of ion is 8, so the structure is described as having (8,8)-coordinationi) Radii are so similar, that this energetically highly favorable coordination is

feasible, with numerous counter-ions adjacent to a given ione) The caesium-chloride structure. The corner lattice points, which are shared by eight

neighboring cells, are surrounded by eight nearest-neighbor lattice points. The anion occupies a cubic holes in a primitive cubic lattice.

f) Caesium-chloride projection

47) Group 2 elements: Beryllium, Magnesium, Calcium, Strontium, Barium, Radiuma) Ca, Sr, Ba, and Ra are known as the alkaline earth metals, but the term is often

applied to the whole of group 2b) All the elements are silvery white metals and the bonding in their compounds is

normally described in terms of the ionic modeli) Some aspects of the chemical properties of beryllium are more like those of a

metalloid with a degree of covalence in its bondingc) Elements are denser, harder, and less reactive than the elements of Group 1, but are

still more reactive than many typical metalsi) Lighter elements, beryllium and magnesium form a number of complexes and

organometallic compounds

(1) Magnesium (Magnesium, Mg)(a) 8th most abundant element in earth’s crust, 3rd in seawater(b) Commercially extracted from seawater and mineral dolomite

(CaCO3•MgCO3)(2) Extracted by electrolysis of their molten chlorides:

(a) Calcium (Calcium, Ca)(i) 5th most abundant in earth crust, 7th in seawater (low solubility of

CaCO3)(ii) Occurs widely in its carbonate as limestone, marble, and chalk, and it

is a major componenet of biominerals (shells and corals)(b) Strontium (Strontium, Sr)(c) Barium (Barium, Ba)

d) Greater mechanical hardness and higher melting points of Group 2 compared with Group 1 indicates increase in the strength of metallic bonding on going from Group 1 to Group 2

48)Beryl (Beryll (Aquamarin), Be3Al2[Si6O18])a) Beryllium aluminum cyclosilicateb) Aquamarine is a blue or turquoise variety

i) Color attributed Fe2+

49) The polymorphs of calcium carbonate: occurs as vast depositories of sedimentary rocks formed from fossilized remains of marine creaturesa) Aragonite (Aragonit, CaCO3)

i) Aragonie is a less abundant polymorph of calcite(1) orthorhombic(2) three crystal forms

ii) less stable than calcite, converts to calcite at 400°C / after enough timeiii) Most bivalve animals secrete aragonite, have aragonite shells

(1) Pearlization due to several layers of aragoniteiv) Synthetic aragonite

(1) used in paper industry: fine texture, whiteness, absorbent properties(2) powdered, in kiln to form lime (CaO), slurried with water to form milk of

lime(a) carbon dioxide bubbled through until aragonite formed

(b) conditions, such as temperature and flow rate of carbon dioxide determine final particle size distribution and crystal size

b) Calcite (Calcit (Doppelspat), CaCo3)i) Most common and most stable formii) hexagonaliii) composes ~4% by mass of earth’s crust

(1) significant part of all three major rock classification types (igneous, sedimentary, metaporphic)(a) Carbonatite in igneous rock(b) limestone is sedimentary form of calcite

(i) limestone metamorphoses to marble from heat/pressure and pure white marble results

iv) Iceland spar (Doppelspat)(1) form of transparent, colorless calcite, originally found in iceland(2) exhibits birefringence (double refraction - division of a ray of light into two

rays when it passes through certain types of material depending on polarization of light; as two beams exit crystal, they’re bent into two different angles of refraction)

50) Gypsum (Gips, CaSO4•2 H2O)a) The most important sulfate is calcium sulfate

i) ‘Permanent hardness’ (not removed by boiling) caused by magnesium and calcium sulfates

b) occurs naturally as gypsum and alabasteri) Gypsum is the dihydrate CaSO4•2 H2O

(1) mined and used as building material(a) Fireproof wallboard: in the event of fire, dihydrate will dehydrate to form

hemihydrate and release water vaporii) Alabaster is a dense, fine-grained form of same molecular composition that

resembles marbleiii) When dihydrate heated above 150°C, loses water to hemihydrate (CaSO4•½H2O)

(1) Plaster of paris(2) Adding of water restores dihydrate and expands

51) Colemanite (Colemanit, CaB3O4(OH)3•H2O)a) Borate mineral found in evaporite deposite

i) Naturally occurring boron compounds are borates52) Fluorite (Fluorspar) (Fluorit (Flussspat), CaF2)

a) The anhydrous halides of the heavier Group 2 elements can be prepared by dehydration of the hydratesi) All the fluorides except BeF2 are sparingly soluble, although the solubiliity

increases slightly down the groupii) As radius of cation increases from Be to Ba, cation coordination number

increases from 4 to 8, with CaF2, SrF2, and BaF2 adopting the fluorite structure(1) Other halides of Group 2 form layer structures, reflecting the increasing

polarizability of the halide structure in which the layers are arranged so that the Cl- ions are cubic close packed

(a) Both MgI2 and CaI2 adopt closely related cadmium-iodide structure in which layers of I- ions are hexagonal close packed

iii) Most important fluoride of group is CaF2, whose mineral form (fluorite or fluorspar) is the only large-scale source of fluorine(1) Anhydrous hydrogen fluoride is prepared by the action of concentrated

sulfuric acid on fluorspar:CaF2(s) + H2SO4(l) CaSO4(s) + 2 HF(l)

b) Primary source of F is calcium fluoride, which has low solubility in water and is often found in sedimentary deposits as fluorite

c) Hydrofluoric acid (HF) is a weak acid, but one of the most toxic and corrosive substancei) readily absorbed through skin and even brief contact can lead to severe burning

and necrosis of the skin and deep tissue and damage to bone decalcification by formation of CaF2 from calcium phosphate

53)Fluorite structure

a) A common AX2 structural type is the fluorite structure, takes its name from naturally occurring mineral fluoritei) The fluorides of larger cations (from Ca to Ba) adopt the (8,4)-coordinate fluorite

structure

ii) Ca2+ ions lie in an expanded ccp array and the F- ions occupy all the tetrahedral holes(1) Cations are close-packed because the F- anions are small

iii) Lattice has (8,4)-coordination, consistent with there being twice as many anions as cations(1) consistent with there being twice as many anions as cations(2) Anions in their tetrahedral holes have four nearest neighbors(3) Cation site surrounded by a cubic array of eight anions(4) Alternative description, is that anions form a nonclose-packed, preimitive

cubic lattice and the cations occupy half the cubic holes in this lattice54)The most important strontium minerals are Celestine and Strontianit

a) Celestine (Coelestin, SrSO4)b) Strontianite (Strontianit, SrCO3)

55) Witherite (Witherit, BaCO3)56) Group 13 elements

a) Elements of Group 13 show a wide variation in abundance in crustal rocks, oceans, and atmospherei) Abundance

(1) Aluminum is the most abundant(a) Aluminum occurs in numerous clays and aluminosilicate minerals, but

the commercially most important mineral is bauxite, a complex mixture of hydrated aluminum hydroxide and aluminum oxide

(2) Low abundance of heavier members of the group is in keeping with the progressive decrease in nuclear stability of elements that follow iron

ii) Whereas elements of the s and d blocks are all metallic, the elements of the p block range, resultin in diverse properties and distinctive trends

iii) Aluminum (Aluminium, Al)iv) Gallium (Gallium, Ga)v) Indium (Indium, In)vi) Thallium (Thallium, Tl)

57) Natural Occurrence of the Elements

58)Use of Aluminuma) The most commercially important element in aluminum; most widely used

nonferrous metalb) Technological uses of aluminum exploit its lightness, resistance to corrosion, and

recyclability59) Corundum (Korund, -Alα 2O3): The corundum structure is adopted by many oxides of

the stoichiochemistry M2O3

a) -Aluminum oxide (the mineral α corundum) adopts a structure that can be modelled as a hexagonal close-packed array of O2- ions with the cations in two-thirds of the octahedral holes

60) Spinel (Spinell, MgAl2O4)a) Spinel itself is MgAl2O4, and oxide spinels, in general, have the formula AB2O4

i) spinel structure consists of a ccp array of O2- ions in which the A cations occupy one-eigth of the tetrahedral holes and the B cations occupy half the octahedral holes

ii) Sometimes denoted A[B2]O4

(1) square brackets denote the cation type (normally the smaller, higher charged ion of A and B) that occupies the octahedral holes

iii) Examples of compounds(1) Many ternary (three parts) oxides with the stoichiometry AB2O4 that contain

a 3d-series metal(2) Some simple binary d-block oxides (in these structures A and B are the same

metal, but in different oxidation states)

b) The observation that many d-metal spinels do not have the normal spinel structure is related to the effect of ligand-field stabilization energies on the site preference of the ions

61) The Group 14 Elementsa) Germanium is a metalloid, and tin and lead are metals (increase in metallic

properties on descending a group is a striking feature of the p block and can be understood in terms of the increasing atomic radius and associated decrease in ionization energy, ie energy required to remove electrons, down the group)i) Because ionization energies of heavier elements are low, metals form cations

increasingly readily down the groupb) As the valence configuration ns2np2 suggests, the +4 oxidation state is dominant in

the compounds of the elementsi) Major exception is lead, most common is 2+

(1) Relative stability of low oxidation state is an example of the inert-pair effect, a striking feature of the heaviest p-block elements

c) All elements of group except lead have at least one solid phase with a diamond structurei) Cubic phase of tin, called grey tin or -tinα ( -Snα ) is unstable at RT, converting

to more stable white tin or -tinβ ( -Snβ )(1) -tinβ contains Sn with six nearest neighbors in a highly distorted octahedral

array

d) Tin (Zinn, Sn)e) Lead (type metal) (Blei (Lettern), Pb)

i) Durch Sb-Zusätze gehärtes Pb nennt man Hartblei. Wichtige Legierungen sind Bleilagermetalle (60-80% Pb, Sb, Sn und etwas Alkali- bzw. Erdalkalimetall) und Letternmetall (70-90% Pb, Sb, und erwas Sn).

62) - α und -tinβ

a) Cubic phase of tin, called grey tin or -tinα ( -Snα ) is unstable at RT, converting to more stable white tin or -tinβ ( -Snβ )i) -tinβ contains Sn with six nearest neighbors in a highly distorted octahedral

arrayb) When white tin cooled to 13.2°C converts to grey tin

63) Galena (Bleiglanz, PbS)a) Lead occurs naturally as galena

64)Lead oxide (PbO, Blei(II)oxid)a) Forms of lead oxide

i) Brown lead(IV) oxide, PbO2

ii) Red and yellow forms of lead(II) oxide, PbO

iii) Mixed oxide, Pb3O4, which contains Pb(IV) and Pb(II)(1) Red lead

b) Relative stability of an oxidation state in which the oxidation number is 2 less than the group oxidation number is and example of the inert pair effecti) recurring theme within p blockii) no simple explanation

(1) often ascribed to the large energy needed to remove the ns2 electrons are the np1 election has been removed

(2) low M-X bond enthalpies for the heavier p-block elements and the decreasing lattice energy as the atomic radii increase down a group

c) The +2 oxide becomes more stable on going down the group from Ge to Pbi) The red form of PbO has the same structure as blue-black SnO with a

stereochemically active lone pairii) Lead also forms mixed oxidation state oxides, best known being red lead, Pb3O4

(1) Contains Pb(IV) in an octahedral environment and Pb(II) in an irregular six-coordinate environment

(2) Assignment of different oxidation numbers to the lead in these two sites is based on shorter PbO distances for the atom identified as Pb(IV)

iii) The maroon form of lead(IV) oxide, PbO2, crystallizes in rutile structure(1) The oxide is a component of the cathode of a lead-acid battery

65) What is a transition element?a) One terms the elements of groups 3-12 as well as the elements Ce to Lu (58-71) and

Th to Lr (90-103) as transition metalsb) They’re located in the periodic table in the region between the metals of Groups 1

and 2 and the above all non-metal elements of groups 13-18i) The term transition element refers exclusively with metals, for which the

expression transition metal is synonymousc) A typical transition element in at least one stable oxidation state has at its disposal

only partially occupied d or f orbitalsi) Besides which the elements of Group 12 also belong to the transition elementsii) One differs between the “outer” transition metals (elements of Groups 3-12, d-

block elements) and “inner” transition metals (partially occupied f orbitals, f-block elements)

66)Subgroupsa) The heavy homologs of Groups VIII (Ru, Os), IX (Rh, Ir), X (Pd, Pt) are called the

Platinum group metals (PGM)i) Reason: they accumulate together in the anode sludge during copper refining

b) The metals of group 11 are called the coinage metals

c) The inner transition metalls are separated into the lanthanoids (elements yttrium as well as lanthan to lutetium) and the actinoids (elements actinium to lawrencium)i) Also called f-block elements

ii) The lanthanoids are also named the rare earth metalsiii) One denotes the inner transition elements without Y, La, Ac as

Lanthanoids/Actinoids (i.e. lanthan-like/actinium-like)67)Electron configuration

a) The (n+1) s and the n d orbitals have nearly equivalent energiesi) With the elements that initially become the (n+1) s shell, then only n d shells are

occupiedb) Electron configuration of the metals

c) Ionization occurs initially from the (n+1) s orbitalsi) Reason: the (n+1) s electrons shield the n d outer-electrons more strongly from

the nuclear charge than other n d electrons(1) Ionization from the (n+1) s orbitals reduce the energy of the n d orbitals from

there more strongly, than ionization from the n d orbitals the remaining n d electrons and the electrons of the (n+1) s shells would be reduced

68)Electron configuration of the ions

a) Please note: Transition metal ions always have (in contrast to the corresponding elements) the electron configuratino (n+1)s0, (n-1)fx, n dy

69)What’s special about the transition metals?a) Variable oxidation state

i) Metals of the main group elements (Groups 1, 2 and 13-17) form in only one of the oxidation states or form in around two units of differing oxidation states stable ions(1) Examples: Na+, Ba2+, Al3+, Tl+, Tl3+, Pb2+, Pb4+

ii) Transition metals, for the most part, form stable ions in several different oxidation states(1) These differ often only by one unit and can be easily converted into one

another

iii) The maximum possible oxidation state corresponds to the group number and has the nd0 configuration(1) These are no longer reached, according to the elements of Group VIII and

with the 3d metalls after Mn (Group VII)b) Stability in low oxidation states

i) in contrary to particular ligands (so-called -acid-ligands, for instance CO, πalkenes), the transition metal ions form stable complexes in small or even negative oxidation states

(1) With metal carbonylates (anionic carbonyl complexes: most neutral metal carbonyl complexes can be reduced to an anionic form known as a metal carbonylate), the oxidation state can reach -4 (for instance, [Mo(CO)4]4-)

c) Colored bondsi) Most of the bonds of the main group metals are colorless or white, whereas

bonds of the transition metals are often intensely colorfulii) Reason: The excitation of electrons within the d or f orbitals or excitation, which

with a shift from the electron density from the metal ions to the ligands (MLCT) or from ligands to metal ions (LMCT)(1) most charge-transfer complexes involve electron transfer between metal

atoms and ligands(a) The charge-transfer bands in transition metal complexes result from shift

of charge density between molecular orbitals (MO) that are predominantly metal in character and those that are predominantly ligand in character

(2) If the transfer occurs from the MO with ligand-like character to the metal-like one, the complex is called a ligand-to-metal charge-transfer (LMCT) complex

(3) If the electronic charge shifts from the MO with metal-like character to the ligand-like one, the complex is called a metal-to-ligand charge-transfer (MLCT) complex(a) Thus, a MLCT results in oxidation of the metal center, whereas a LMCT

results in the reduction of the metal centerd) Unpaired electrons and magnetism

i) partially occupied d/f orbitals and energy differences between individual groups of d and f orbitals often lead to the presence of one or several unpaired electrons (how many maximum?)(1) And that’s why they’re paramagnetic

(a) Form of magnetism, whereby the paramagnetic material is only attracted in the presence of an externally applied magnetic field

ii) Nonetheless, they frequently show no radical-type reactivity at all and are stablee) The formation of complexes of variable coordination number and coordination

geometryi) This is a general characteristic of all metal ions and will be come across also with

the main group elementsf) Physical characteristics

i) transition metals often have at their disposal:(1) High heat capacity and conductivity (Cu!)(2) High ductility (solid material’s ability to deform under tensile stress, often

measured according to its ability to be stretched into a wire)(3) to some extent, high melting and boiling points (for instance, Smp. ~3380°C,

Sdp ca. ~5500°C)(a) High densities (heavy homologs of the later transition metals)

(4) Partly high resistance of metal to oxygen, acids, and bases (precious metals)g) Oxidation states

i) Under the (formal) oxidation states (oxidation numbers), one understands this charge, which remains on the central atom, if all of these atom binding groups are conceptually separated in their corresponding best Lewis representation(1) Outer charges must be, in addition, taken into account

70)The effective nuclear charge Zeff

a) Electrons in a multiple electron system are found in a Potential, which is given through the attraction between electron and nucleus Vatt = - 1/(4πε0) • ze2/rek and the repulsion between the electrons vpot = 1/(4πε0) • Σij e2/rij

i) rij is the distance of the electron i and of electron j from one anotherii) One simply accepts, that electrons of one through n, l describe sets of quantum

numbers through the nuclear proximity or in orbitals with the same n.iii) l residing electrons are “shielded from” the positive nuclear charge

b) The effective potential of an electron behaves therefore according to Vatt = - 1/(4πε0) • ze2/rek, whereby zeff is the so-called effective nuclear chargei) This is defined as zeff = z - Σi≠j σi

(1)σi is the atomic screening constant of the electron i(2) electron j is being observed

c) Electrons of the same quantum numbers n and l are mutually shielded only weakly from the nuclear charge (screening constant = 0.35)σi) Because of the large number of electrons with the same n and l, this effect for

transition metals is especially strongly reflectedii) Within the n d or n f rows, the effective nuclear charge zeff increases greatly with

increasing atomic number(1) Two important consequences:

(a) Within a n d or n f row, the atomic radii decrease with increasing atomic number (in other words, from left to right)(i) Within the d-block elements with the 5d metals, this effect is

especially strongly reflected

(ii) The reason is the occupation of the 4f shell with lower mutual shielding as well1. Through which, the zeff increases with increasing atomic number

even more greatly

2. Atoms (ions) of metals, which come after the lanthanoids, are much smaller than expecteda. Most have practically the same atomic/ionic radius as their

corresponding lighter homologsb. This feature is called the lanthanide/actinide contraction

d) Ionization energiesi) Two consequences of the increasing of zeff:

(1) The first ionization energies tend to increase with transition metals, that are within a row and within a group, with increasing atomic number

(2) Observation: the heavier elements of a group are more difficult to oxidize and develop stabler bonds(a) With higher atomic number, the contribution of the d orbitals in the

development of chemical bonds increases(i) Example: Pt(PR3)2/H2 and Pd(PR3)2/H2

71)Area of stabler oxidation statesa) The “earlier” transition metals (Groups 3-5) and the “later” transition metals

(Groups 9-12) have comparatively less stable oxidation statesi) The “earlier” transition metals have only fewer d electrons at their disposal,

which are given away or can participate in bondingii) The “later” transition metals have too few unoccupied orbitals or too high

ionization energies to reach higher valences72)Redox potential

a) Within the d row, the 4d elements have (especially in complexes), as a rule, the highest oxidation potentials

i) Example:

ii) Example: 73)Precious metals (Edelmetalle)

a) Noble Characteri) Kupfer (Copper, Cu)ii) Rubidium (Rubidium, Ru)iii) Rhodium (Rhodium, Rh)iv) Palladium (Palladium, Pd)v) Silber (Silver, Ag)vi) Rhenium (Rhenium, Re)vii)Osmium (Osmium, Os)viii) Iridium (Iridium, Ir)ix) Platin (Platinum, Pt)x) Gold (Gold, Au)xi) Quecksilber (Mercury, Hg)

b) With many redox reactions, the redox potential depends on the pH valuei) An example of which are the reactions of metals with acids and water

(1) All metals with negative potential, in other words all metals in the electrochemical series above hydrogen, can with it give away electrons to the H3O ions and develop oxygen(a) One calls these metals the base/common metals (unedle Metalle)

(2) Metals with positive potential, which are located below hydrogen (Cu, Ag, Au) can’t be dissolved in acid by H-develoopment and are, for instance, in HCl insoluble(a) These metals are the precious metals (edle Metalle)

(3) From this, with water, all metals should be able to react during the development of hydrogen, whose potential is more negative than -0,41 V

(a) Base metals have a negative standard potential, precious metals a positive one(i) Only base metals are dissolved in acids by hydrogen evolution

74) Foundations of coordination chemistrya) Definition of Complexes / Coordination bondsb) Historyc) Nomenclatured) Ligandse) Coordination numbers / Coordination polyhedraf) Isomerismg) Molecular symmetry

75)Jørgensen - Werner

a) Alfred Werner, U. Zürich, “father of complex chemistry”, 1913 Nobel prizei) Conductivity measurements

76)Coordination theory according to Alfred Werner (1898)

77)Coordination polyhedra and the number of isomers

78)IUPAC Nomenclaturea) Identify central atom(s)

i) For complicated structures, the name is easier to form if more central atoms are chosen

b) Identify ligandsc) Name ligands

i) Anionic ligands require special endingsd) Specify coordination mode for each ligand:

(1) specify donor atom(s)(2) specify central atom(s)

ii) The conversion is generally applicableκiii) Note that is used when contiguous atoms are coordinatedη

e) Order ligands and central atom(s)i) Ligand names are ordered alphabeticallyii) Central atom names are ordered according to their position in table

f) Identify coordination geometry and select polyhedral symboli) Most structures will deviate from ideal polyhedra, so choose closest

g) Describe relative configurationi) CIP (Cahn-Ingold-Prelog) priority used

(1) Compare the atomic number (Z) of the atoms directly attached to the stereocenter; the group having the atom of higher atomic number receives higher priority

(2) If there is a tie, we must consider the atoms at distance 2 from the stereocenter—as a list is made for each group of the atoms bonded to the one directly attached to the stereocenter. Each list is arranged in order of decreasing atomic number. Then the lists are compared atom by atom; at the earliest difference, the group containing the atom of higher atomic number receives higher priority.

(3) If there is still a tie, each atom in each of the two lists is replaced with a sub-list of the other atoms bonded to it (at distance 3 from the stereocenter), the sub-lists are arranged in decreasing order of atomic number, and the entire structure is again compared atom by atom. This process is repeated, each time with atoms one bond farther from the stereocenter, until the tie is broken.

h) Determine absolute configuration79)Important terms relating to complexes

a) The metal-bound Lewis bases in complexes are called Ligandsi) (Chelatligand): Binding a ligand over more than one Lewis base center on the

same metal ionii) (Brückenligand): Binding on two or more different metal ions

80)Names for anionic and neutral ligands

81) Denticity (Zähnigkeit) of Ligandsa) Monodentate ligands

i) Donor atoms:(1) Nitrogen(2) Oxygen(3) Sulfur(4) Phosphorus(5) Halogens(6) Bonds can also function as Lewis bases

ii) Polydentate ligands(1) Chelate ligands

(a) Greater thermodynamic stability(b) Entropy effect

(2) Bridging ligands(a) can bridge two or more metal centers

82)Important terms relating to complexesa) Nature of ligands

i) A complex that contains only one type of ligand are called homolepticii) Complexes with multiple different ligands present are heteroleptic

b) The oxidation state / oxidation numberi) This formalism doesn’t always lead to sensible bond splitting

(1) Example hydride ligands

(2) Observation: the oxidation numbers are based on a formalism and are not physically measurable sizes!(a) Nevertheless there are measurable qualities that are correlatable with

the oxidation state (nuclear magnetic resonance shift of metal nuclei, Mössbauer spectrums)

c) The d electron numbersi) The d electron numbers correspond to the difference between the group number

and the oxidation state of the metal ion(1) Example: [NiCl2(en)]

(a) Ni has the oxidation level +II and is aGroup 10 element: the d e- # is 8(2) Example: [Mn(NO3)2(SO4)2]2-

(a) Manganese has the oxidation state +IV and is located in Group 7; the number of electrons is 3

ii) To remember: metal ions in complexes always have the electron configuration (n+1)s0, ndm

d) The valence electron numbers (VENs)i) how many valence electrons does a complex need?

(1) The 18 valence electron rules of Sidgwick (1927) say that thermodynamically stable transition metal complexes exist when the metal atom has over 18 valence electrons at their disposal and so the electron configuration attains the periodic system of the next inert/noble gas

ii) Are there also complexes with higher or lower valence electron numbers?(1) In complexes, the valence electron numbers range from 10 to 20

(a) Examples of complexes with lower VENs

83)Polydentate ligandsa) Examples of bidentate ligands

b) Examples of tri-, tetra-, and hexadentate ligands

84)Crown ether (Kronether): Potassium-[18]-crown-6 (Kalium-[18]-krone-6)a) The space-filling model shows, that the potassium ion fits perfectly in the crown

85)Porphyrin Synthesis: Template effect

86)Polydentate ligands as indicatorsa) Murexid metal complex: the metal complex is yellow, the free ligand is violet

b) Eriochrome black T complex: the metal complex is red, the free ligand is blue

87)Coordination polyhedra and relationships to radii

88)Coordination polyhedra

89)Coordination Number 4 (Quadratic Planar)

90)Coordination Number 4 (tetrahedral)

91)Coordination number 5 (trigonal bipyramidal and quadratic pyramidal)a) Quadratic pyramidal

i) Example [Ni(CN)5]3-

b) Trigonal bipyramidali) Ligand repulsionii) [Fe(CO)5] shows a signal only in 13C-NMR spectroscopyiii) fluctuating molecule (Berry pseudorotation)

c) easily convert to one anotherd) Berry pseudorotatione) Turnstyle mechanism

92)Berry Pseudorotation

a) A berry pseudorotation in which (a) a trigonal-bipyramidal Fe(CO)5 complex distorts into (b) a square-pyramidal isomer and then (c) becomes trigonal bipyramidal again, but with the two initially axial ligands now equatoriali) (a) = 180°, = 120°α βii) (b) = α βiii) (c) = 120°, = 180°α β

93)Turnstyle mechanism

94)Trigonal prisma) C3v symmetryb) Example: [W(CH3)6], [Re(CH3)6]

95)Coordination number 7a) Pentagonal bipyramid

i) Example: [V(CN)7]4-

b) Overcapped trigonal prismi) Example: K2[NbF7]

c) Overcapped octahedra

96)Coordination number 8a) Quadratic antiprism, Archimedes prism

i) D4d symmetryii) Example: H4W(CN)8•6H2O

b) Trigonal dodecahedroni) D2d symmetryii) Example: K4Mo(CN)8•2H2O

c) Cube97)Triple overcapped trigonal prism

a) Example: [ReH9]2-

98)The most important coordination polyhedra

a) Conspicuous: Majority of (and most important ones of) the coordination polyhedra are assembled exclusively from triangular surfaces (=deltahedron)

b) Why are the just discussed coordination polyhedra so especially stable?i) In the main groups chemistry, the VSEPR- (or Gillespie/Nyholm-) rules produce

a simple connection between the composition and the structure of a compound(1) The basis is the minimization of the repulsive force (repulsion) between

electron pairs of the central atom (binding or free)ii) The model of Kepert is a transmission of this simple concept of transition metal

complexes(1) By doing this, one accepts that the central atom is found in the center of a

sphere(2) The ligand donor atoms are arranged on the surface of a sphere around the

central atom (in other words the ligand donor atom distances are fixed on the sphere radius)

iii) The ligands are electrostatically mutually repelled(1) The donor atoms take these positions on the sphere surface, which

minimizes the ligand-ligand repulsion(2) Deltahedra are exactly this polyhedron, for which the ligand-ligand repulsion

is minimal

1) Isomerism

2) Bond Isomerism: Typical ligands

3) Bond Isomerism: Nitrito ligand

4) Bond Isomerism: CN liganda) Berliner Blau (prussian blue)b) Prussian blue

5) Geometric Isomerism - Coordination Number 4a) Tetrahedral

i) No geometric isomersb) Quadratic-planar

i) cis and trans isomers

6) Geometric Isomerism - Coordination Number 5

a) Quadratic-pyramidal

b) Trigonal-bipyramidal

7) Geometric Isomerism - Coordination Number 6a) Octahedral

b) Trigonal-prismatic

8) cis-trans Isomerism (geometric isomerism)a) Quadratic-planar

b) Octahedral

9) Coordination theory according to Alfred Werner (1898)

10)And with three different ligands?11)[MA2B4]

12)[MA3B3]

13)[MA2B2C2]

14)Optical isomerism and chirality

15)Optical activity

16)Optical isomerism and chirality - coordination number 6a) Absolute configurations of M(L-L)3 complexes.

i) is used to indicate clockwise rotation of the helix and to indicate Δ Λanticlockwise rotation

b) [Co(en)3]3+ : Image and mirror image

17)Recognizing chirality

a) If a complex has either a mirror plane or center of inversion, it cannot be chirali) If we look at the schematic complexes in (74: [Cr(edta)]-), (75: [Ru(en)3]2+), and

(76: [Pt(dien)Cl]+), we can see that neither (74) nor (75) has a mirror plane or a center of inversion, so both are chiral

ii) (76) has a plane of symmetry and hence is achiral

18)Bond models for complexesa) Valence bond theoryb) Crystal field and ligand field theoryc) Molecular orbital theory

19)Bond models for coordination compoundsa) Borderline: covalent

i) Lewis acid base conceptii) 18-electron ruleiii) VB theoryiv) MO theory

b) Borderline: electrostatici) ionic modelii) crystal field theoryiii) ligand field theory

20)Hybridizationa) d

21)d-Orbitals (473-)a) There are two models of electronic structure of d-metal complexes:

i) crystal-field theory

(1) ligand lone pair modelled as a point negative charge (or as partial negative charge of an electric dipole) that repels electrons in the d orbitals of the central metal ion

(2) Concentrates on resulting splitting of the d orbitals into groups with different energies, and uses splitting to explain properties of complexes

ii) ligand-field theory

22)Octahedral ligand fielda) In the presence of an octahedral crystal field, d orbitals are split into a lower-

energy triply degenerate set (t2g) and a higher-energy doubly degenerate set (eg) separated by an energy Δp

; the ligand-field splitting parameter increases along a spectrochemical series of ligands and varies with the identity and charge of the metal atom.i) in CFT model of octahedral complex, six point negative charges representing

ligands are placed in octahedral array around central metal ion(1) These charges (referred to from now on as ligands) interact strongly with

central atom(a) Stability of complex stems from attractive interaction between opposite

chargesii) Electrons in dz^2 and dx^2-y^2 orbitals are concentrated close to the ligands, along

the axes, whereas electrons in dxy, dyz, and dzx orbitals are concentrated in regions that lie between the ligands(1) former repelled more strongly by the negative charge on the ligands than the

latter and lie at a higher energy

23)Splitting/breakdown in the octahedral fielda) Group theory shows that

i) two eg orbitals have the same energy (although not readily apparent from drawings)

ii) the three t2g orbitals also have the same energyiii) This simple model leads to an energy-level diagram in which the three

degenerate t2g orbitals lie below the two degenerate eg orbitals(1) Separation of the two sets of orbitals is called the ligand-field splitting

parameter ΔO (where O subscript signifies octahedral crystal field)

b)

c)

i)

24)Electromagnetic Spectrum

25)UV/visible spectrum of [Ti(H2O)6]3+

26)UV/visible spectrum of [Cr(H2O)6]3+

a) Three transitions/transfers in the electron spectrum!

27)Ligand-Field stabilization energiesa) The ground-state configuration of a complex reflects the relative values of the

ligand-field splitting parameter and the pairing energy. For 3dn species with n = 4-7, high-spin and low-spin complexes occur in the weak-field and strong-field cases, respectively. Complexes of 4d- and 5d-series metals are typically low-spin.

b) d orbitals in a complex don’t all have same energy, thus ground-state electron configuration of a complex is no longer obviousi) To predict, use the following d-orbital energy diagram as a basis for the building-

up principle

(1) i.e. we identify the lowest energy configuration subject to the Pauli exclusion principle (a maximum of two electrons in an orbital) and (if more than one degenerate orbital is available) to the requirement that electrons first occupy separate orbitals and do so with parallel spins

(a)c) First we consider complexes formed by the 3d-series elements

i) in an actahedral comples, first three d electrons of a 3dn

complex occupy separate t2g nonbonding orbitals, and do so with parallel spins(1) for example, Ti2+ and V2+ have electron configurations 3d2

and 3d3 respectively(a) The d electrons occupy the lower t2g orbitals as shown

in (1) and (2) respectively(b) The energy of a t2g orbital relative to the barycentre

(energy level that corresponds to the hypothetical spherically symmetrical environment (in which negative charge due to the ligands evenly distributed over a sphere instead of being localized at 6 points) defines the barycentre of the array of levels, with two eg orbitals lying at 3/5ΔO above the barycentre and three t2g orbitals lying at 2/5ΔO below it) of an octahedral ion is -0.4ΔO and the complexes are stabilized by 2•(0.4ΔO) = 0.8ΔO (for Ti2+) and 3•(0.4ΔO) = 1.2ΔO (for V2+)(i) Additional stability relative to barycentre is called

the ligand-field stabilization energy (LFSE)

28)High-spin (HS) and low-spin (LS) complexesa) When alternating configurations are possible, species with the smaller number of

parallel electron spins is called a low-spin complex and greater number is a high-spin complexi) An octahedral 3d4 (as well as 3d5, 3d6, 3d7) complex is likely

(1) low-spin if ligand field strong(2) high-spin if weak

29)Ligand-field stabilization energy (LFSE) for d0-d3

30)LFSE for d4-d5

a) P = pairing energy

31)LFSE for d6-d7

32)LFSE for d8-d10

33)Tetrahedral ligand field

34)Factors determining the size of ΔO

a) The Oxidation state of the metal ioni) ΔO increases with increasing oxidation state of the metal ion

b) The type of metal ioni) ΔO increases within a group from above 3d elements to below 5d elements

(1) Consequence: 4d and 5d elements are almost always in the low-spin configuration

c) The type of ligand (see below)

35)Spectrochemical series: the type of ligand

36)Fe(III) complexes

37)Co(III) complexes

38)Co(II)/Co(III) complexes - Redox potentials

39)Ion radii for M2+/M3+ - The effect of the LFSE

40)Hydration enthalpies for M2+ ions

41)Ligand SALCs and metal orbitals

42)MO diagram for ML6 complexesa) only -bondσ

b) -outward bond and -back bondσ π

43)Consequences from the MO examinationa) complexes with weak donor ligandsσ

i) t2g is non-bonding, eg* is weakly anti-bondingii) both can be, however don’t have to be, occupiediii) high-spin complexes, 12-22 VE

b) complexes with strong donor ligandsσi) t2g is non-bonding, eg* is strongly anti-bonding

ii) Occupation of the eg* is unfavorableiii) low-spin complexes, ≤18 VE

c) complexes with acceptor ligandsπi) t2g is bonding, must be occupiedii) eg* is anti-bonding, occupation is unfavorableiii) 18 VE rule must be satisfiediv) most metals in a low oxidation state

d) complexes with donor ligandsπi) t2g and eg* are anti-bondingii) Occupation of t2g and eg* is unfavorableiii) lowest possible d electron numberiv) most of the early transition metals or metals in high oxidation states

44) Traits of complexesa) Jahn-Teller effectb) Magnetismc) Complex stabilityd) Reactions of coordination compoundse) Trans effect

45)Review: Ligand field splitting

46)Jahn-Teller Theorema) If the electrical ground state of a non-linear molecule is degenerate, then the

molecule will be distorted, such that the degeneration is offset and a lower energy state and symmetry is adopted

47)Octahedral - tetragonal distortion

48)Jahn-Teller effect

49)Types of magnetisma) Paramagnetism

b) Ferromagnetism

c) Antiferromagnetism

50)Exchange pathsa) Antiferromagnetic coupling

b) Ferromagnetic coupling

51)Substitution reactions: on quadratic-planar complexes

a) IA mechanismb) The reaction rate constant is dependent on the nucleophile of the receiving ligand Y

i) Y = PPh3 > CN- > I- > N3- > NH3 > Cl- >> CH3OH

52)Stabilization of the transition state

a) -acceptor ligands stabilize the transition stateπb) Electron density is withdrawn by the metal

53)Rare Earth Metals (Seltenerdmetalle)a) Scandiumb) Yttriumc) Lanthanoids

54)Ion radii of the ions of Ln3+ (lanthanides with 3+ oxidation state)

55)Electron configurations of the lanthanoids

56)Electron configurations of the actinoids

57)Group 4 Elementsa) Titan (Titanium, Ti)b) Zirconium (Zirconium, Zr)c) Hafnium (Hafnium, Hf)

58)Rutile lattice (TiO2)

59)Compounds ABX3 with Perowskite structurea) Oxides ABO3

i) A = Ca, Sr, Ba; B = Ti, Zr, Hf, Sn, Ce, Tc(1) Examples: CaTiO3, BaCeO3

ii) A = Rare earth metals; B = Al, Sc, V, Cr, Mn, Fe, Co, Ga(1) Example: LaMnO3

b) Fluoride KBF3

i) B = Mg, Cr, Mn, Fe, Co, Ni, Cu, Zn

c) Oxide fluorides ANbO2Fi) A = Li, Na, K

d) Chlorides (CsBCl3) and Bromides (CsBBr3)i) B = Cd, Hg, Au, Ga

(1) Examples: CsAuCl3, CsGaCl3

e) Sulfidesi) ATIS3

(1) A = Sr, Baii) AZrS3

(1) A = Ca, Sr, Ba60)Ziegler-Natta Catalyst

a) Nobel prize for chemistry (1963)

61)UV/visible spectrum of [Ti(H2O)6]3+

62)Group 5 Elementsa) Vanadium (Vanadium, V)b) Niob (Niobium, Nb)c) Tantal (Tantalum, Ta)

63)Vanadyl sulphate (VOSO4•5H2O)64)Vanadium(V) oxide (V2O5)65)Oxoanions (Oxyanions) of Vanadium (in their dependence on concentration and pH

value)

66)Halides of V, Nb, and Ta

67)Group 6 Elementsa) Chrom (Chromium, Cr)b) Molybdän (Molybdenum, Mo)c) Wolfram (Tungsten, W)

68)Molybdänit (Molybdenite, MoS2)

69)Oxoanions of Cr subject to pH value

70)Cr Oxides

71) Mo and W Oxides

72)Cr Halides

73)Mo and W Halides

74)Metal-metal multiple bondsa) M-M double bond

i) M-M - and -bondsσ π

b) M-M triple bondi) M-M - and 2 -bondsσ π

c) M-M quadruple bondi) M-M -, 2 -, and -bondsσ π δ

75)Group 7 Elementsa) Mangan (Manganese, Mn)b) Technetium (Technetium, Tc)c) Rhenium (Rhenium, Re)

76)KMnO4

77)Mn Oxidesa) Pyrolusite (Braunstein) (MnO2)

78)Tc and Re Oxides

a) Rhenium(VI) oxide (ReO3)

79)Mn Halides

80)Tc and Re Halides

81)Re-Re quadruple bond

a) The ReCl4 groups are located ecliptically with regard to one another

82)Isopolymolybdate (1)

83) Isopolymolybdate (2)

84)Keggin Structure

85)Group 8 Elementsa) Eisen (Iron, Fe)b) Ruthenium (Ruthenium, Ru)c) Osmium (Osmium, Os)

86)The most important iron oresa) Magnetite (Fe3O4)b) Siderite/Eisenspat (Iron Ore) (FeCO3)c) Hämatit (Hematite)/Roteisenstein (red hematite)/Blutstein (bloodstone) (Fe2O3)

87)Pyrite (FeS2)

88)Blast furnace process

89)Gicht des Hochofens (Blast furnace top of a blast furnace)90)Fe Halides and Oxides

a) FeCl3 Hexahydrate

91)Ru and Os Halides and Oxidesa) OsO4

92)Gelbes Blutlaugensalz (yellow prussiate of potash, K4[FeII(CN)6])

93)Rotes Blutlaugensalz (red prussiate of potash, K3[FeIII(CN)6])

94)Berliner Blau (Prussian blue)

95)Group 9 Elementsa) Cobalt (Cobalt, Co)b) Rhodium (Rhodium, Rh)c) Iridium (Iridium, Ir)

96)Uses of Co

97)Co Halides and Oxidesa) CoCl2•6 H2O = [CoCl2(H2O)4]•2 H2O

98)Rh and Ir Halides and Oxides

99)cis and trans isomerism of [CoCl2(NH3)4]Cla) Violeo (cis) (Greek: green)

b) Praseo (trans) (Latin: violet)

1) Group 10 Elementsa) Nickel (Nickel, Ni)b) Palladium (Palladium, Pd)c) Platin (Platinum, Pt)

2) Important nickel oresa) Rotnickelkies (niccolite, NiAs)b) Gelbnickelkies (millerite, NiS)c) Weißnickelkies (chloanthite, NiAs2-3)

3) Ni Halides and oxidesa) NiCl2•6 H2O = [NiCl2(H2O)4]•2 H2O

1) Pd and Pt Halides and Oxides

2) Group 11 Elementsa) Kupfer (Copper, Cu)b) Silber (Silver, Ag)c) Gold (Gold, Au)

3) Cu Halogens, Oxides, and Sulfidesa) Cu2O

4) Ag Halides, Oxides, and Sulfides

5) Au Halides, Oxides, and Sulfides

6) CuSO4•5 H2O = [Cu(H2O)4]SO4•H2O

7) Group 12 Elementsa) Zink (Zinc, Zn)b) Cadmium (Cadmium, Cd)c) Quecksilber (Mercury, Hg)

8) Zinnober (Sinnople, HgS)

9) Zn and Cd Halogens, Oxides, and Sulfidesa) CdS

10)Zincblende and Wurtzite (ZnS)

11)Hg Halides, Oxides, and Sulfidesa) HgO

12)Special ligandsa) Hydride ligands and Dihydrogen Ligandsb) Oxide ligands and Dioxygen ligandsc) Dinitrogen and Nitrosyl Ligandsd) Cyanido Ligandse) Carbonyl Ligands

13)Dihydrogen Complexes

14)Dioxygen Ligand

15)Metal coordination of O2

16)Dioxygen complexes: Orbitalsa) Hin-Bindung (end-on)

b) Rück-Bindung (end-on)

c) Hin-Bindung (side-on)

d) Rück-Bindung (side-on)

17)Vaskas Complex

a)

18)O2 end-on bridging

a)

19)Dioxygen ligands vs. dinitrogen ligands: MO Diagram

20)Dinitrogen ligand

21)Dinitrogen complexes

a)

b)

c)

22)Nitrosyl Ligand

23)Nitrogen monoxide

24)Nitrosyl Ligand

25)Metal coordination of NOa) [CpFe( -NO)]μ 2

b) [(PPh3)ClIr( -NO)( -dppm)PdCl]PFμ μ 6

c) [Fe(CN)5(NO)]2-

d) [Co(NH3)5(NO)]Cl2

26)Nitrosyl ligand: When is it linear and when is it bent?a) Analyzes using 15N-NMR, IRb) Ligand NO+

i) corresponding CO is linearii) 3 e- donor

c) Ligand NO-

i) corresponding O2 is bentii) 1 e- donor

27)Nitrosyl Complexes: MO diagram

28)Cyanide ligand

29)Cyanide/Isocyanide liganda) Isocyanide

b) Cyanide

30)Coordination Polymers

31)Berliner Blau (Prussian blue)

32)Carbonyl ligand

33) interactionπa) Effects the LF splitting Δ

34)Qualitative MO diagram of CO

35)Mononuclear CO complex, 18 VE rule

36)Homoleptic, mononuclear neutral CO complexa) Why don’t Ti, Hf, and Zr; Pd and Pt form stable, homoleptic, neutral carbonyl

complexes?

37)Homoleptic, multinuclear neutral CO complexes

38)Dinuclear CO complexesa) Mn2(CO)10

b) Fe2(CO)9

c) Co2(CO)8

39)Tri- and tetranuclear CO complexesa) Os3(CO)12

b) Ir4(CO)12

c) Fe3(CO)12

40)Synthesis of metal carbonyls

41)Bond relationships

42)Oscillation spectroscopy and structure

43)Spectroscopy dates of [M(CO)6]n+/-

44)Synthesis of carbonyl metalates