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1!CHEM112 LRSVDS Transition Metals part 1!

Transition Metals

•! Occupy the d-block of periodic table

•! Have d-electrons in valence shell

Characteristics:

1.! More than one oxidation state

2.! Many compounds are colored

3.! Interesting magnetic properties

4.! Form Metal complexes or Coordination compounds

5.! Transition metals play important roles in biological systems and modern technology.

2!

Electron Configurations and Oxidation States Many transition metals form compounds that have fun colors!

– colors are due to oxidation state and electron configuration...more specifics about that later!

3!

Periodic trends

There are periodic trends in the transition metals, but they

are often complex (product of several factors, some working in opposite

directions – e.g. combining the effects of increasing nuclear

charge with the presence of nonbonding d electrons)

Lanthanide contraction – similarity in

size, behavior, properties of 4d

and 5d transition elements

We won’t worry about details of

periodic trends in the

transition metals or the exact

reasons for them


Electron Configuration: Order of Orbital Filling


VALENCE ELECTRON CONFIGURATIONS

Orbital Filling in First Row: Sc ! Zn

[Ar] = 3s23p6 K [Ar] 4s1

Ca [Ar] 4s2 Sc [Ar] 3d14s2

Ti [Ar] 3d24s2

. .

. .

. . Zn [Ar] 3d104s2

*Note: 4s is filled before 3d, but when metal is oxidized, 4s electrons are lost before 3d.

Ti [Ar]3d24s2

Ti2+ [Ar]3d24s0

Ti3+ [Ar]3d14s0 Ti4+ [Ar]3d04s0

Ti5+ does not exist!


Group 3B – 7B:

For Sc, Ti, V, Cr, Mn: highest oxidation states common Highest oxidation states = number of valence (4s + 3d) electrons.

Sc [Ar]3d14s2 Sc3+ [Ar] (maximum oxidation state is +3)

Mn [Ar]3d54s2 Mn7+ [Ar] (maximum oxidation state is +7)

Group 8B, 1B, 2B: The maximum oxidation state becomes increasingly unstable and uncommon.

Sc3+ Sc2O3 is a stable oxide

Mn7+ Exists but is easily reduced (MnO4! strong oxidizing agent)

Fe8+ Does not exist; unstable

Transition Metal Oxidation States


Metal Complexes or

Coordination Compounds

Transition metal ions are Lewis acids "

they accept electron pairs

Ligands are Lewis bases "

molecules or ions which donate electron pairs


Ligands: Donate Lone Pairs of Electrons

Anionic Ligands

F!, Cl!, Br!, CN!, SCN!, NO2!, EDTA4!

Neutral Ligands

NH3, H2O, CO, CH3OH, en

Mono-dentate Ligands: Only ONE donor atom is bound to the metal (single tooth to hold onto metal d orbital)

Examples; NH3, H-O-H , CH3-O-H

Bi-dentate Ligands: TWO donor atoms are bound to the metal (has 2 teeth to hold onto metal d orbitals).

Example; H2N-CH2-CH2-NH2 (en or ethylenediamine)

Poly-dentate Ligands Have more than one functional group with lone pairs (many teeth!)

Example; EDTA (ethylenediaminetetraacetic acid)


Coordination Compounds

Coordination number:

number of donor atoms attached to the metal.

Chelates: “chele or Claw”

ligands possessing two or more donor atoms.!

[Cu(H2O)2(NH3)2]2+ [Cu(H2O)2(en)2]

2+

en =!


TRANSITION METAL COMPLEXES

[Cu(NH3)4)]SO4 SO42- + [Cu(NH3)4)]

2+

Charge on the complex:

Coordination #:

Oxidation state of the metal:

K2[Ni(CN)4)] 2 K+ + [Ni(CN)4]2-

Charge on the complex:

Coordination #:

Oxidation state of the metal:

11!

Werner’s Theory

•! This approach correctly

predicts there would be two

forms of CoCl3 · 4 NH3.

–! The formula would be written [Co

(NH3)4Cl2]Cl.

–! One of the two forms has the two

chlorines next to each other.

–! The other has the chlorines

opposite each other.


Tetrahedral

e.g. [Zn(NH3)4]2+

Square Planar

e.g. [Ni(CN)4]2-

e.g. [PtCl2(NH3)2]

Geometry of Transition Metal Complexes:

Coordination Number of 4

Pt

Cl

Cl NH3

NH3

13!

Geometries of Transition Metal Complexes Geometry for Coordination # = 5

•! Trigonal Bipyramidal

[Fe(CO)5]

[Re(SCH2C6H4OCH3-p)3(PPh3)2] ReL3(PR3)2!


e.g. [CoF6]3-

e.g. [Co(en)3]3+!

Six Coordinate Complexes:

Octahedral (six vertices, eight FACES)

Co

F

F F

F

F

F

Co

N

N N

N

N

N


Poly-dentate Ligands: Form Metal CHELATES!

ethylenediamine (“en”)!

= ethylenediaminetetraacetate ion!


NCH2CH

2N: :

CH2COH

CH2COH

O

O

HOCCH2

HOCCH2

O

O

EDTA

IMPORTANT CHELATING LIGANDS


Uses of Chelating Agents

•! Used to “sequester” or “seize” metal ions

•! Used in detergents to remove trace amounts of dissolved metals: Na5P3O10

•! Complex trace metal ions that catalyze food decomposition: EDTA

•! Used in poison control: EDTA

•! Used in shampoo and cleaning products to remove trace metals from hard water (Ca2+ and Mg2+): EDTA


IMPORTANT CHELATING LIGANDS

Porphine forms metal complexes called Phorphyrins

Fe!

!"#$%#&'()

+ Fe + protein ! !


Chemistry and Life!

!"#$%$&"'##()!

*+%%,-"%$.+!

http://fr.academic.ru/dic.nsf/frwiki/29449!

http://en.wikipedia.org/wiki/File:Ferrichrome.png!


Important Chelating Agents

Chelate # of Coordination Charge

Sites

Ethylenediamine

Porphine

EDTA4-

Oxalate (C2O42-)

Carbonate (CO32-)

21!

Metal Complexes and Isomers

22!

Structural Isomers

•! Coordination Sphere Isomers –! If a ligand (like the NO2

group at the bottom of the complex) can bind to the metal with one or another atom as the donor atom

•! Linkage Isomers –! differ in what ligands are

bonded to the metal and what is outside the coordination sphere

Three isomers of CrCl3(H2O)6 are:

violet [Cr(H2O) 6]Cl3

green [Cr(H2O) 5Cl]Cl2·H2O

green [Cr(H2O) 4Cl2]Cl·H2O

23!

Stereoisomers

•! Geometric Isomers •! Optical Isomers


METAL COMPLEX STABILITY

Cu(OH2)42+ + 4NH3 [Cu(NH3)4]

2+ + 4H2O(l)

Cu2+(aq) + 4NH3 [Cu(NH3)4]2+ + 4H2O(l)

[H2O] = constant

[Ag(NH3)2]+ 1.7 x 107

[Cu(NH3)4]2+ 5.0 x 1012

[Ag(S2O3)2]3! 2.9 x 1013

[Ag(CN)2] ! 1.0 x 1021

[Cu(CN)4]2! 1.0 x 1025

Formation Constant: Kf is very large!

Kf = !


THE CHELATE EFFECT Chelating ligands form exceptionally stable metal complexes.

[Ni(H2O)6]2+ + 6NH3 [Ni(NH3)6]

2+ + 6H2O

Kf = 4x108

[Ni(H2O)6]2+ + 3en [Ni(en)3]

2+ + 6H2O

Kf = 2x1018

*DUE TO:

1) PROBABILITY

2) ENTROPY EFFECTS!

Cd+2

NH2CH3

NH2CH3

Cd+2

H2N

NH2

Probability Effect:!


Entropy and the Chelate Effect 1) Cd2+ + 4CH3NH2 [Cd(CH3NH2)4]

2+

2) Cd2+ + 2en [Cd(en)2]2+!

Why is #S° so much larger?

[Cd(H2O)4]2+ + 4CH3NH2 [Cd(CH3NH2)4]

2+ + 4H2O

[Cd(H2O)4]2+ + 2en [Cd(en)2]

2+ + 4H2O

Ligand #H°(kJ) #S°(J/K) #G°

1 methyl amine -37.2kJ

2 en -60.7kJ

27!

Diamagnetic:

unaffected by a magnetic field

Paramagnetic:

influenced by a magnetic field

The magnetic properties depend on the number of unpaired electrons

Na+

Mn2+

Ti2+

Co3+

CHEM112 LRSVDS Transition Metals part 1!

Magnetic Properties of Transition Metals


Magnetic Behavior

Diamagnetic!

Paramagnetic!

Ferromagnetic!

# unpaired e"! prior alignment alignment in magnetic field !attracted to magnetic field !


The presence of metal-ligand bonding electrons raises the energy

of metal d orbitals due to electrostatic repulsion.

E

d orbitals in free metal ion (all degenerate)

!

d orbitals in uniform, “spherical” field of negative charge; all orbitals raised in

energy equally

CRYSTAL FIELD THEORY


Octahedral Metal Complexes; Effect of metal-

ligand bonding electrons on metal d orbitals

Which metal d orbitals are most affected by the metal-

ligand bonding electrons?!

Metal ion in Octahedral charge Field


CRYSTAL FIELD SPLITTING

Crystal Field Energy Splitting of d orbitals in octahedral ligand field!

d orbitals in octahedral field of negative charge

!

d orbitals in uniform, “spherical” field of negative charge

!

e set (dz2, dx2–y2)

t2 set (dxy, dxz, dyz)!

#$ “delta octahedral”

E

#!o= Crystal field splitting energy

!o depends on:

1.

2.

3.


Spin Pairing Energy P = spin pairing energy

The energy required to place two electrons of opposite

spin in the same orbital

*magnitude of P is independent of the ligands

If # is large (P < #) " Low Spin Complex

If # is small (P > #) " High Spin Complex

For a transition metal with 5 d electrons;!

First 3 electrons fill lower E orbitals!


CN!

CO

NO2!

en

NH3

H2O

Oxalate

OH-

F!

SCN!

Cl!

Br-

I!

Strong field!

ligands!

Weak field !

ligands!

Increasing # %

absorbs!

Cl! < F! < H2O < NH3 < en < NO2! < CN!

SPECTROCHEMICAL SERIES:

Ability of L to increase the energy gap

observed!


MAGNETIC PROPERTIES

Which diagram corresponds to CoF63- and

which corresponds to Co(CN)63- and why? !

E

High spin

Paramagnetic!

Low spin (spin-paired)

Diamagnetic

35!

Practice Problems

CHEM112 LRSVDS Transition Metals part 1!

1. Ammonia is a strong field ligand. Is [Mn(NH3)6]3+ high spin or low spin?

2. The oxalate complex [Co(C2O4)3]4& has 3 unpaired electrons. Is it high

spin or low spin?


OPTICAL PROPERTIES; COLOR OF COMPLEXES

Observed color is related to the amount of

energy required to promote an electron.

Compare # to energy absorbed.


[Ni(H2O)6]2+ + 6 NH3 ! [Ni(NH3)6]

2+ + 6 H2O

Color depends on identity of the ligands

OPTICAL PROPERTIES OF TRANSITION METALS


COLOR ABSORPTION BY METAL COMPLEXES

When light of a certain wavelength is absorbed by a complex, the complex will appear the complementary color of the

wavelength absorbed

Observed Color:

1)

OR

2)


Visible Absorption Spectra:!

What color is Ti(H2O)63+?!


Color of Metal Complexes

1. Which of these complexes absorbs light at a shorter wavelength?

2. Which complex has the larger #o?

3.! Which is a weaker field ligand; water or thiocyanate?

4. What color will TiO2 be? What color will ZnO be?

[Fe(H2O)6]3+!

SCN"!

[Fe(SCN)(H2O)5]2+!


Reduction Potential of Metal Complexes

Stability of Transition Metal Complexes depends on

reduction potential of the metal complex.

Which is easier to reduce, the metal ion or the

complex?

Ag+(aq) + e- ! Ag(s) E°1/2= +0.80V

[Ag(CN)2]-(aq) + e- !Ag(s)+ 2CN-(aq) E°1/2 = -0.31V


USES OF TRANSITION METALS Titanium

Vanadium

Chromium

Manganese

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