1

complex: used by chemists for compounds that consist of (several) other compounds that can exist separately

Coordination compounds (complex)

2

ligands

central atom

coordination sphere (inner)

acceptor

donor

Lewis base

Lewis acid

No. of donors exceeds the

value of oxidation number

usually (poly-nuclear) ions + counterions

Coordination compounds (complex)

3

Some of further basic concepts

Coordination number: number of donor atoms coordinated in the inner sphere

Ligands: monodentate – a single donor atom (H2O, CN-, F- … )

polydentate – their geometry enables to occupy (bi-, tri- ...) more than a single coordination position several donor atoms (chelate agents) (e.g. ethylendiamin H2N-CH2-CH2-NH2)

4

chelate complexes

Ethylendiamin (en)

EDTA

bridging ligands

Ethylendiamintetraacetate(4-)

5

Coordination compounds: bonds/structure

Alfred Werner, Swiss, 1866-1919, Nobel Prize 1913

Showed that transition metals create complexes with square, tertrahedral, octahedral structure

e.g. cis-[PtCl2(NH3)2] trans-[PtCl2(NH3)2] diammin-dichloridoplatinum(II) complex

geometrical isomers

cis-

trans-

6

Typical space structures of complexes

Trigonálny dodekaéder Trojnásobne zastrešená

trigonálna prizma

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Geometrical isomerism for octahedral structures

cis- trans- mer- fac-

Info: Optical isomerism: mirror image – enatiomers

chirality, chiral molecules

8

Valence bond theory with hybrid AO in most cases enables explanation of the structure

coord. No. form of coord. sphere examples 2 – SP linear [CuCl2]- [Ag(S2O3)2]3-

4 – SP3 [Co(NCS)4]2- [NiCl4]2- D3S tetrahedron [BF3(NH3)]

4 – DSP2 [Mn(H2O)4]2+ [PdCl4]2- SP2D square [Pd(NH3)4] 2+ Ni(CN)4]2-

6 – D2SP3 [Fe(H2O)6]2+ SP3D2 octahedron [Fe(CN)6]3- [FeF6]3- [PdCl6]2-

Coordination compounds: bonds/structure

paramagnetic [NiCl4]2- unpaired electrons

Ni(II) -[NiCl4]2–

Ni2+

sp3

9

28Ni 3d 4s 4p

High-spin complex

diamagnetic [Ni(CN)4]2- paired electrons

Ni(II) -[Ni(CN)4]2–

Ni 3d 4s 4p

Ni2+

dsp2

10

Ni2+ valence

Low-spin complex

11

metal-ligand bond is weaker than „usual“ covalent b.

??????

Some complexes use inner „d“ orbitals others use outer „d“ orbitals

Transition metal complexes use to be intensively colored

MO theory

simplified approximations

Coordination compounds: bonds/structure

12

Crystal field theory Central atom in electrostatic field

of (ionic) ligands (as point charges) (electrostatic theory of ligand field)

Splitting of „d“ levels: octahedral complex

d

Ene

rgy Δ

Ligand field

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Why the transition metal complexes are colored?

d-d transitions

eg

t2g

splitting of „d“ levels: tetrahedral complex

d

Ene

rgy Δ

t2

e

Ene

rgy

octahedral

complex is violet

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low and high-spin complexes [Fe(CN)6]3- [FeF6]3-

Fe3+

[Fe(CN)6]3-

5d Fe0

4s

Δ

[FeF6]3- E

nerg

y

Δ

t2g

eg

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Relative ligand field strengths

High-spin complexes Low-spin complexes

Spectrochemical series

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Oxidation-reduction (redox) reactions: change of oxidation state

-IV, -III ... 0, I, II ... VIII oxidation

reduction

x BAox1 + y DCox2 x EAox1+m + y FCox2-n

Reducing agent (electron donor)

oxidizing agent (electron acceptor)

electron balance: donated e– = accepted e –

x.m = y.n

Example:

partial redox equations

Br2 2Br + 10e– 0 V

Cl + 2e– Cl I -I

Br2 + HCl O + H2O HBr O3 + HCl 0 I V -I

Br2 2Br + 10e – 0 V

5 Cl + 10e – 5 Cl I -I

balancing

5 2 5

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Oxidation-reduction (redox) reactions: disproportionation reactions

(dismutation reactions)

Cl2 + H2O HCl O + HCl 0 I -I

x BAox1 + y CAox1 x EAox1+m + y FAox1-n

(x+y) Aox1 x Aox1+m + y Aox1-n

3KCl O3 + KCl O3 3KCl O4 + KCl V V -I VII

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Introduction to electrochemistry: electrode, electrode potential

metal/H2O metal ions

hydrated metal ions

+ + + + +

+ + + + + +

equilibrium: M(s) = Mz+ + z.e-

Vel Vr Vr -Vel

in the solution according to cM z+

E = E0 + ln cM z+ R.T

z.F

Nernst equation – electrode potential

Faraday const. = NA.e

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Galvanic cells

Half cell Half cell

Galvanic cell

Interrupted circuit EMV = E2-E1 electromotoric voltage

In a closed circuit: electronic flow, ion flow (electric current) electric

work

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Standard electrode potentials

absolute n/a

25°C E = E0 + log cM z+ 0.059

z

Standard electrode potential: cM = 1.00 mol L-1 z+

Reference electrode: H3O+/H2

Standard hydrogen electrode: E0 =0.0 V

1)   platinized platinum electrode 2)   H2 blow (1 atm)

3)   Solution of an acid [H3O+] = 1 mol L-1 4)   hydroseal for prevention of O2 intervention

5)   junction to the second half cell

H3O+ + e– = ½ H2 + H2O

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Standard electrode potential: measurement

Standard electrode potentials

reducing ability Noble m

etals non-noble

Metals with lower E0

reduce the cations of metals with higher E0

cathode

-1.18 eV

-0.13 eV

Mn+Pb2+ Mn2++Pb

Oxidation/reduction (redox) potentials:

redox system + indifferent (Pt) electrode

Ox + z e– = Red

reducing ability

E = E0

+ log 0.059

z [Ox] [Red]

Nernst-Peters equation

Electrode

Example: H2O2 +2H++Sn2+ Sn4++2H2O

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Reverse process: Electrolysis

+

E0=1.36 V

Na+ (aq) + e- = Na(s) E0= –2.07 V

2H2O(l) + 2e- → H2(g) + 2OH-(aq) Cl2(g) + 2e- → 2Cl-(aq)

E0= -0.83 V

preferable process

in progress: 2Cl-(aq) → Cl2(g) + 2e-

E = -0.41 V [OH-]=10-7

1.7-2.2 V

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Electrolysis from molten salts:

Cathode: Na+ + e– → Na(l) E° = –2.71 v

Anode: Cl– → ½ Cl2(g) + e– E° = –1.36 v

together: Na+ + Cl– → Na(l) + ½ Cl2(g) E° = –4.1 v

Na

Al