CHEM1010 Transitio Metal Lec Notes

THE UNIVERSITY OF QUEENSLAND AUSTRALIA

CHEM1010 Transition Metal Chemistry School of Chemistry and Molecular Biosciences



TRANSITION METAL CHEMISTRY

Dr Philip Sharpe CHEMISTRY

PHONE 3365 3900 Room 205 Chemistry

EMAIL [email protected]

CHEM1010



Associate Lecturer at UQ since 2010. I’m the academic in charge of the first year Chemistry laboratory. I also coordinate CHEM1021 and teach second year biological inorganic chemistry in CHEM2052. My research interests are in the area of chemistry education and coordination chemistry applied to medicine.



This section of the lecture course deals with

material covered in Chapter 13

of Blackman “Chemistry”.

Note: First edition cannot be used for inorganic nomenclature section.

These notes are incomplete and

require you to add material.

Figures in these lectures notes are taken from Blackman, 2nd Edition, and are used with permission from the publishers and distributors John Wiley







Why is this stuff

(transition metal chemistry) important?

Note: This first section on why transition metals are important is not examinable.



Ruby is an aluminium oxide (Al2O3) crystal in which some of the aluminium atoms have been replaced with chromium atoms. Chromium gives ruby its

characteristic red colour.



Chromium is responsible for the lasing behavior of the crystal. Chromium atoms absorb green and blue light and emit or reflect only red light.

http://www.llnl.gov/nif/library/aboutlasers/how.html



Metal Ions in Biology

Worldwide prevalence of anaemia 1993-2005, WHO Global Database on Anaemia



Essential Transition Metals Biological Function Vanadium Essential for rats and chickens. No

specific role known in humans. Manganese Needed for several enzymes, e.g. Mn

superoxide dismutase, glutamine synthetase.

Iron Hemoglobin, myoglobin, other enzymes

Cobalt Vitamin B12

Nickel Urease (plants), hydrogenase (bacteria)

Copper Found in several enzymes. Cu-Zn superoxide dismutase detoxifies free radicals.

Zinc Bound to insulin, around 10% of all proteins contain Zn.

Molybdenum (only essential 2nd row transition metal)

Component of redox enzymes e.g. sulfite oxidase.



hemoglobin (Hb) - a tetrameric protein O2 uptake in the lungs and

transport in the blood stream. 65% of the iron present in a human.

Other iron proteins – involved in DNA synthesis, drug metabolism, energy pathways and many other functions.



The organic core of hemoglobin

Porphyrin Heme (Porphyrin

+ iron)



http://www.chemistry.wustl.edu/~edudev/LabTutorials/Hemoglobin/changemovie.html

A heme site in hemoglobin

http://www.chemistry.wustl.edu/~edudev/LabTutorials/Hemoglobin/changemovie.html�





DINITROGEN - thermodynamically stable, so reduction requires a large amount of energy. N2 + 10H+ + 8e- → 2NH4

+ + H2

In the Haber Process, N2 and H2 gases are reacted over an Fe3+ catalyst in which Al2O3 and K2O are used as promoters. The reaction is carried out under conditions of 250 atm, 450-500 °C; resulting in a yield of 10-20%:

N2(g) + 3H2(g) → 2NH3(g) ΔHθ = -92.4 kJ mol-1

The Haber process produces 100 million tons of nitrogen fertilizer per year, mostly in the form of anhydrous ammonia, ammonium nitrate and urea. 1% of the world's annual energy supply is consumed in the Haber process. That fertilizer is responsible for sustaining 50% of the Earth's population.



How Nature does it:-

Rhizobium bacteria, which live in legume root nodules contain nitrogenase protein.

MoFe7S8



Science 334, 974 (2011) DOI: 10.1126/science.1206445 Kyle M. Lancaster, et al. X-ray Emission Spectroscopy Evidences a Central Carbon in the Nitrogenase Iron-Molybdenum Cofactor CARBON!!!!!!!!!!!!!!! Starts out as a methyl group in S-adenosylmethionine – the process is still being investigated. J. A. Wiig, Y. Hu, C. C. Lee, M. W. Ribbe, Science 337, 1672 (2012).



Metal ions in Medicine • cis-[Pt(NH3)2Cl2] – cisplatin; an anticancer drug • Radionucleotides – 64Cu for tumour imaging • Antiarthritic drugs – gold complexes Metal ions in Industry • Catalysts e.g. - Optically active metal complexes as agents for synthesis of chiral pharmaceuticals. - catalytic converters in cars for removal of nitrogen oxides from exhaust (Rh, Pt, Pd).



We are at a wonderful time for chemistry. It is, I believe, in the position of physics in the 1910s, just before quantum mechanics made the world impossibly strange, or biology in the 1950s, just before the double helix obliterated the old biology. George M. Whitesides Priestley Medal Address, 2007.







The Aufbau Principle – allows us to write electron configurations

Blackman p 138, 546

http://z.about.com/d/chemistry/1/0/j/g/econfiguration.jpg�



Transition metal atoms:- • characterised by d valence orbitals (e.g. 3d, 4d, 5d) • neutral atoms generally have valence electron configurations of (n+1)s2nd(x-2) (x is the group number of the metal in the periodic table, n is the principal quantum number) e.g. Vanadium, group number of 5, electron configuration [Ar]4s23d3. IUPAC definition: A transition metal is "an element whose atom has an incomplete d sub-shell, or which can give rise to cations with an incomplete d sub-shell". http://goldbook.iupac.org/T06456.html

Blackman p 545-546

Zn, Cd, Hg (Group 2B elements) are d-block metals, but we can think of them as honorary transition metals because of their similar chemistry.



It is all about the d-orbitals & d-electrons!!!

The five 3d orbitals means a maximum of 10 d electrons. Two

orbitals lie on axes, three orbitals lie

between axes.





for example, element Ni (Z = 28); group 10 1s2 2s2 2p6 3s2 3p6 4s2 3d8

Remember from Module 1: • Quantum numbers • orbitals.

•All 3d orbitals are degenerate (equal in energy) in the elements and ions in the gas phase.

http://z.about.com/d/chemistry/1/0/j/g/econfiguration.jpg�



1s2 2s2 2p6 3s2 3p6 4s2 3d1 Sc 21 Ti 22 V 23 Cr 24 Mn 25 Fe 26 Co 27 Ni 28 Cu 29 Zn 30

1s2 2s2 2p6 3s2 3p6 4s2 3d2

1s2 2s2 2p6 3s2 3p6 4s2 3d3

1s2 2s2 2p6 3s2 3p6 4s2 3d10

1s2 2s2 2p6 3s2 3p6 4s1 3d10

1s2 2s2 2p6 3s2 3p6 4s2 3d8

1s2 2s2 2p6 3s2 3p6 4s2 3d7

1s2 2s2 2p6 3s2 3p6 4s2 3d6

1s2 2s2 2p6 3s2 3p6 4s2 3d5

1s2 2s2 2p6 3s2 3p6 4s1 3d5

Cr and Cu don’t follow the regular pattern. The 4s and 3d orbitals in atoms are close in energy, so the lowest energy configuration has the 4s orbital half-filled and the 3d orbitals half-filled (Cr) or completely filled (Cu).



Oxidation States • Transition metals are typically found in various oxidised forms i.e. they can lose one or more electrons to form cations.

Fe → Fe2+ + 2e- Fe → Fe3+ + 3e- • The chemical and physical properties are extremely dependent on the oxidation state (colour, reactivity, structure). • The electrons that are lost come from the 4s orbital first, before the 3d orbitals. Remember that these orbitals are close in energy in atoms in the gas phase, and they swap relative energy levels in ions compared to the atoms. IMPORTANT!!!

Blackman, p 546



Table 13.1, p. 546 Blackman 2nd ed. • For ions in the same oxidation state, the number of d electrons increases going across the row from left to right. •For the same element, the number of d electrons decreases as oxidation state increases. •Elements in the middle have the greatest number of possible oxidation states.

= common oxidation states

p. 546 Blackman 2nd ed.



The higher oxidation states are limited by:- • The effective nuclear charge – it becomes energetically unfavourable to remove more electrons. • The number of d electrons. • The intermediate oxidation states are often unstable due to disproportionation reactions leading to more stable reduced and oxidized forms. 2Mn3+ → Mn2+ + Mn4+

• Low oxidation states are unstable due to spontaneous oxidation by air. 2Cr2+ + O2 + 2H+ → 2Cr3+ + H2O2



An Aside:- “Ionization Energy” The ionization energy of an element is defined as the energy change when an electron is removed from an atom in the gas phase. For an element M the ionization energy (I1) is the energy change in the process M(g) → M+(g) + e (I1) (kJ mol-1) And M+(g) → M2+(g) + e (I2) (kJ mol-1) M2+(g) → M3+(g) + e (I3) (kJ mol-1)



0

500

1000

1500

2000

2500

3000

3500

4000

Sc Ti V Cr Mn Fe Co Ni Cu Zn

Ionisation Energy

(kJ mol-1)

Element

1st Ionisation energy 2nd Ionisation energy 3rd Ionisation energy

Gas phase ionisation energies for the first row TM elements



In transition metal chemistry we are often interested in knowing the number of d electrons (not least for exam questions). Easiest way is

to count them using the Periodic table.

The number of d electrons of transition metal ions The electrons that are lost from first-row TM atoms to form ions come from the 4s orbital first, before the 3d electrons are lost. (5s and 4d in the case of second row transition metals)

1 2 3 4 5 6 7 8 9 10

For a divalent (2+) transition metal ion, the two 4s electrons are lost and then we just count along the row starting from scandium, the first transition metal.

(For Cr2+ and Cu2+ ions, the one 4s electron will be lost and one 3d electron.)



Ni2+ : Ni Z=28, so a Ni2+ ion has 28 - 2 = 26 electrons Last noble gas was Ar (Z=18). No. of 3d electrons = 26 -18 = 8 electrons 1s2 2s2 2p6 3s2 3p6 3d8

Really important!!!! 6/18 TM questions in last year’s exam required knowledge of the number of d-electrons.

Alternative method: Number of d-electrons = Number of electrons in ion – number electrons in last noble gas element.



21 22 23 24 25 26 27 28 29 30

Sc Ti

V Cr Mn

Fe Co Ni Cu Zn

2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+ 2+

d1 d2 d3 d4 d5 d6 d7 d8 d9 d10

Atomic Number Element Ionic Charge d electron Config.

Divalent first row metal ions



What about TM ions with oxidation states other than II? Example: Co3+ This will have one less d-electron than Co2+ (or 3 less electrons than Co)

1 2 3 4 5 6 7 8 9 10

Co2+ has 7 d electrons, so Co3+ will have 6 d electrons. Alternately, Co (Z=27), so Co3+ has 27 - 3 = 24 electrons. Ar (Z=18) was the last noble gas, so number of d electrons = 24 -18 = 6 d electrons.



Which two commonly occurring first-row transition metal ions have FIVE 3d electrons?

A. Co3+ and Ni2+ B. Fe2+ and Co3+

C. V5+ and Cr6+

D. Mn2+ and Fe3+

E. Co4+ and Ni5+

Blackman worked example 13.1, p 546





Coordination Chemistry d-block metal ions are Lewis acids and they form compounds with Lewis bases. Lewis bases donate electron pairs. Lewis acids (e.g. d-block metal ions) accept electron pairs. The Lewis base donates an electron pair to the metal ion. The Lewis bases that are bonded to the metal ion in this way are called ligands. ([From Latin ligandus, to bind) The compound formed by a metal ions and its ligands is called a complex.



Coordinate Bonding Transition metal ions form special type of bonds with other molecules/ions called coordinate bonds

- May be described as a Lewis acid-base reaction. Lewis acid: electron pair acceptor – the metal ion. Lewis base: electron pair donor – the ligand.

Blackman p 552



Consider BF3 (B not a transition metal, example only)

B 1s2 2s2 2p1

3 valence electrons

B

F

F

F

BF3 has properties consistent with a Lewis structure with 3 B-F single bonds

6 electrons around the B.



B

F

F

FN

H

H

H B

F

F

FN

H

H

H

Both electrons in the B-N bond are contributed by N.

This is a coordinate covalent bond

N B Lone empty Pair orbital

N B

Coordinate bonds are normal two electron bonds.



LIGANDS

Metal ion

Ligand

Complex or

complex ion Any atom within

a molecule or ion bearing a lone pair of electrons can potentially coordinate to a metal ion. NH3 can use the lone pair of electrons on the N to coordinate.

The coordinated atom is known as the donor atom.



•A metal ion can have several donor atoms coordinated at the same time.

Blackman p 547-551



A monodentate ligand – NH3

Blackman p 547-551

LIGANDS

Other common monodentate ligands are F-, Cl-, Br-, I-, H2O.

Dentate = having teeth

Octahedron An octahedral ligand arrangement



Denticity

The number of donor groups from a given ligand attached to the same central atom.

bidentate = interacts with a metal through two donor atoms tridentate = interacts with a metal through three donor atoms tetradentate = interacts with a metal through four donor atoms pentadentate = interacts with a metal through five donor atoms hexadentate = interacts with a metal through six donor atoms

Very much a connotation of holding on, e.g. “teeth”



Chelate (Chelation (from Ancient Greek χηλή, chelè, meaning claw) Ring

Blackman p 547-551

A bidentate ligand – 1,2-diaminoethane

or “ethylenediamine”

5-membered chelate ring 6-membered chelate ring

A bidentate ligand – pentane-2,4-dionato or acetylacetonato

or “acac”

http://en.wikipedia.org/wiki/Greek_language�



compare

monodentate

chelate

(bidentate)

en = 1,2-diaminoethane

Blackman p 553

http://images.google.com.au/imgres?imgurl=http://www.pendotech.com/images/mad_scientist.gif&imgrefurl=http://www.pendotech.com/pendotech_home/mad_scientist.htm&h=450&w=338&sz=120&tbnid=tmwyxqLIOWUcBM:&tbnh=127&tbnw=95&prev=/images?q=mad+scientist&um=1&start=2&sa=X&oi=images&ct=image&cd=2�



Name of ion or molecule

Formula Name of coordinated ligand

Abbreviation Example

ethylenediamine 1,2-diaminoethane (ethylenediamine)

en [Ni(en)3]2+

Propane-1,3-diamine

Propane-1,3-diamine

pn [Co(pn)3]3+

2,2’-bipyridine 2,2’-bipyridine bipy [Ru(bipy)3]2+

1,10-phenanthroline

1,10-phenanthroline phen [Ru(phen)3]2+

Pentane-2,4-dionate ion

Pentane-2,4-dionato acac [Co(acac)3]

oxalate oxalato ox2- K3[Cr(ox)3]

1,4,7-triazaheptane

1,4,7-triazaheptane dien [Co(dien)2]3+



a polydentate ligand - EDTA

Fig 13.6, p. 548 Blackman [Co(EDTA)]-

Ethylenediaminetetraacetic acid



Blackman p 549-550

Multidentate ligands can have the same or mixed donor atoms (e.g. EDTA).



Ligands can be linear, branched, cyclic or bicyclic.

[Cu2(H2O)2 (µ -OOCCH3)4] Figure 13.7, Blackman 2nd ed.

Ligands can potentially bridge two or more metal centres.



Many different types of ligands!!!

Blackman p 551

C5H5- ligand, an example

of an organometallic ligand where C atoms are bound to the metal. The complex has the nickname “ferrocene”.



Ligand design is important in -the development of new materials for the capture and storage of CO2 and hydrogen. - new medical diagnostic agents. - new drugs to target metalloenzymes. - improved catalysts.

Science 329 (5990) 424-428 (2010) DOI: 10.1126/science.1192160



Dr Vamvounis – consultation “drop-in” session on organic chemistry; podium Friday May 24 from 10-12 Transition metal drop-in Friday May 24 from 2:10-4. Quiz 3 available now!



Transition metal ions are the most common Lewis acids (electron pair acceptors) Combinations of transition metals and ligands result in a complex (a coordination compound) with a defined geometry.

Blackman p 554



Ligand

Metal ion Ionic charge

Complex or complex

ion

Coordination number = count the number of donor atoms attached to the metal ion.

Coordination number = 6, as there are 6 N atoms coordinated.

Note: There will be a counterion or ions balancing any charge but outside the complex ion. This is indicated by them being written outside the square brackets.

[Co(NH3)6]3+



Oxidation state of the metal ion is calculated by knowing • Coordination number (= 6) • The charges on the ligand (0) • The type and number of counter ion (3 x Cl-, a monoanion) • The overall charges must cancel: x + 3 x -1 = 0, so must be Co(III) = Co3+.

http://en.wikipedia.org/wiki/Hexamminecobalt%28III%29_chloride



Summary of Metal Complexes: Contain coordinate covalent bonds. Lewis acid-base adducts. Composition: central metal ion+ligands+counter ion/s (if needed) Called a complex ion if charged. Coloured, often with unusual magnetic properties e.g. [Cu(NH3)4]SO4

Central metal

ion

ligands (inside the square brackets)

counterion (outside the square brackets)



Consider CoCl3.nNH3 with n = 4, 5 or 6 4 different compounds are possible. How can we have these compounds with various compositions? Dissolve in water, add AgNO3. Some Cl– precipitates as AgCl immediately, some more slowly; this difference indicates the number of Cl– which are ionic (fast) and covalent (slow) Composition formula colour Number Cl– precipitated CoCl3.6NH3 [Co(NH3)6]Cl3 yellow 3 CoCl3.5NH3 [Co(NH3)5Cl]Cl2 purple 2 CoCl3.4NH3 [Co(NH3)4Cl2]Cl green 1 CoCl3.4NH3 [Co(NH3)4Cl2]Cl violet 1

ionic covalent



Coordination number 2: collinear. Collinear complexes - common in the case of heavy metal cations of d10 electron configuration. [Au(CN)2]− , formed during the extraction of gold from its ore, and [Ag(NH3)2]+ , formed when AgCl dissolves in ammonia solution. Coordination number 3: trigonal planar - quite rare; these complexes are found in instances in which ligands are large and steric repulsions are dominant, for example, [Pt{P(phenyl)3 }3]. Coordination numbers 7, 8, and 9: various Further types of coordination geometry exist for large cations, especially those of the 3+ lanthanide cations, for example, [La(edta)(H2O)3]-, [NbF7] 2− , [Mo(CN)8]4− , and [ReH9]2− .

Coordination Geometries Blackman p. 553



Coordination number 4: tetrahedral - common for d10 metal ion species such as Zn2+ and Ga3+ , and d7 species, such as Co2+ . Examples: [ZnCl2(pyridine)2]; [GaCl4]− ; and [CoCl2(4–methylpyridine)2]. Square planar – [Cu(NH3)4]2+ , also common for Ni(II), Pt(II), Ir(I), Rh(I) d8 systems. There is more space to fit ligands around a tetrahedral metal centre, so it is especially favoured for small metal ions with bulky ligands, as the ligands in a square planar complex will be closer together.

Note: Mistake p. 554 d10 metal ion should be Cu(I) not Cu(III). Cu(III) is d8.



Coordination number 5: trigonal-bipyramidal and square-pyramidal - common in the case of complexes of metal ions of coordination number 5. Examples: trigonal-bipyramidal complexes:- [CuCl5]3−, [Fe(CO)3(PF3)2], and [Mn(CO)4NO]. Examples: square-pyramidal complexes:[Ni(CN)5]3− and [Cu(hfacac)2(OH2)] . (hfacac = hexafluoroacetylacetonate)

Blackman p. 554



Coordination number 6: Octahedral Donor atoms at either (approximately) 90° (cis) or 180° (trans) to each other.



Coordination compounds may

exhibit isomerism, just like organic compounds (only better). (Do you

remember this from the early organic

chemistry lectures?)



Isomerism in Coordination Compounds

Blackman p 555-558

Structural isomers (different bonds, same

molecular formula)

Linkage isomers (different

atoms in the same ligand bound to the

metal ion)

Coordination sphere isomers

Ionisation isomers

Stereoisomers (same bonds, different

arrangements)

Hydrate isomers

Isomers (same formula, different

properties)

Geometric isomers

Optical isomers

cis and trans fac and mer Coordination isomers



• Interchange of counter ions and ligands. • Physical properties are completely different.

Blackman p 555-558

Ionisation isomers



Results from the different numbers of water molecules that can be coordinated to a metal ion. A hydrate is formed when a crystal contains loosely held water molecules that are NOT coordinated to the metal ion. [Cr(OH2)6]Cl3 purple [Cr(OH2)5Cl]Cl2·H2O blue-green [Cr(OH2)4Cl2]Cl·2H2O green

Blackman p 555-558

Hydrate isomers



- This form of isomerism results when ligands are exchanged between a complex cation and a complex anion of the same coordination compound.

Blackman p 555-558

[Co(NH3)6][Cr(CN)6] [Cr(NH3)6][Co(CN)6]

Coordination isomers



Requires two potential donors on the one ligand: e.g. NO2

-, CN-, SCN-. These ligands are called ambidentate, as they have more than one potential donor atom.

Blackman p 555-558 Linkage isomers

(which atom in a ligand is bound to the

metal ion)



cis-[Co(NH3)4Cl2]+ trans-[Co(NH3)4Cl2]+

Two ligands at 90° Two ligands at 180°

Blackman p 555-558

Geometric isomers

cis and trans

Stereoisomers – isomers with the same order of attachment of atoms, but a different orientation of their atoms in space.

Stereoisomers (same bonds, different

arrangements)



Geometric Isomerism

Blackman p 555-558

Cisplatin is a chemotherapy drug that is given as a treatment for some types of cancer. It is most commonly used to treat testicular, bladder, lung, gullet, stomach and ovarian cancers. Trans isomer is inactive.

cis-[Pt(NH3)2Cl2] trans-[Pt(NH3)2Cl2]



Facial (fac)

Blackman p 557

Look at the set of three ammine (NH3) ligands. All three NH3 ligands are: - on the same face of the octahedron. - cis (90°) to each other.

This is the facial isomer.

Geometric isomers fac and mer



http://www.plingfactory.de/Science/Atlas/Kennkarten%20Algen/Diatomeen/Source/Meridion%20sp..html

Meridional (mer)

Look at the set of three ammine (NH3) ligands. - All three NH3 ligands are coplanar with the metal ion and each other. - Two of the NH3 ligands are trans (180°) to each other. One NH3 ligand is cis to the two others.

Meridion species of diatom (algae)



mer-[Co(NH3)3Cl3]

fac-[Co(NH3)3Cl3]

http://upload.wikimedia.org/wikipedia/commons/6/69/Fac-trichlorotriamminecobalt(III).png�

http://upload.wikimedia.org/wikipedia/commons/5/54/Mer-trichlorotriamminecobalt(III).png�



Facial isomer – 3 ligands all cis to each other.

Meridional isomer – 2 ligands trans to each other, one ligand cis to the two others.



Stereoisomerism Stereoisomers – isomers with the same order of attachment of atoms, but a different orientation of their atoms in space. Enantiomers – stereoisomers that are nonsuperimposable mirror images of each other.

Figure 17.10, p. 749 Blackman 2nd ed.

Optical isomers





[Co(ox)3]3-

Are these isomers of each other?



Optical isomers - enantiomers

Are they mirror images? Yes!!



Are they superimposable? No!!

∴They are enantiomers, as they are non-superimposable mirror images.

If you can rotate two structures so that they are identical

then they are not isomers. If you have to break bonds to make them identical, they are isomers.



Consider the following complex ion. Can this complex ion exist as optical isomers?

Yes it can

These are mirror images which cannot be superimposed.



Consider the following complex ion: are cis and trans isomers possible?

1. Cis and trans isomers possible 2. Cis and trans isomers not possible



cis (two optical isomers)

trans No optical isomers



Can the following molecule exist in enantiomeric forms?

1. Yes it can exist as optical isomers 2. No it cannot exist as optical isomers



Non-superimposable mirror images – therefore they are optical isomers (enantiomers).



RULES FOR NAMING METAL COMPLEXES

(i) Cations are named before anions, with a space between the ions. This is the same rule that applies to other ionic compounds e.g. sodium chloride.

(ii) The names of anionic ligands always end in the suffix “-o”. Ligands with names ending in ide, ite or ate have this suffix changed to ido, ito and ato, respectively. e.g. a chloride ion when bound to a metal is a chlorido ligand. a carbonate ion when bound to a metal is a carbonato

ligand

Look on the Blackboard page!!!!

Blackman p 558-560 2nd edition only



(iii) A neutral ligand is given the same name as the neutral molecule e.g. 1,2-diaminoethane Except: H2O Water becomes aqua NH3 Ammonia becomes ammine (Note: double m!!!) CO Carbon monoxide becomes carbonyl

(iv) When there is more than one of a particular ligand, their number is specified by the prefixes di = 2, tri = 3, tetra =

4, penta = 5, hexa = 6. If the name of the ligand already incorporates one of these

prefixes (e.g. 1,2-diaminoethane), enclose the ligand name in brackets and use the

following prefixes instead bis = 2, tris = 3, tetrakis = 4. e.g. [Cu(en)2]Cl2 is bis(1,2-diaminoethane)copper(II) chloride

NOT di(1,2-diaminoethane)copper(II) chloride.



(v) In the name of the complex, the ligands are named first, in alphabetical order without regard to charge, followed by the name of the metal. Ignore the prefix in

front if any. e.g. cis-[Pt(NH3)2ClBr] cis-diamminebromidochloridoplatinum(II) A B C



(vi) Negative (anionic) complex ions always end in the suffix –ate. The suffix is appended to the English name of the metal atom in most cases (or Latin stem for

elements where the element symbol comes from Latin). If the name of the metal ends in ium or ese the ending is dropped and replaced by –ate.

metal name in an anionic complex

iron (Fe) ferrate

copper (Cu) cuprate

nickel nickelate

zinc zincate

cobalt cobaltate

chromium chromate

manganese manganate

Note: mistake p. 560 Blackman [CoCl2(NH3)4]- tetraamminedichloridocobaltate(III) ion



(vii) The oxidation state of the metal is included in the name in parentheses in capital Roman numerals. Having this information included means that there is no need to specify the number of counter ions. K4[Fe(CN)6] potassium hexacyanidoferrate(II) NOT tetrapotassium hexacyanidoferrate(II) K3[Fe(CN)6] potassium hexacyanidoferrate(III)



Name of ion or molecule Formula Name of coordinated

ligand

example

water H2O Aqua [Cr(H2O)6]2+

ammonia NH3 Ammine [Co(NH3)6]3+

Carbon monoxide CO Carbonyl [Ni(CO)4]

Fluoride F- Fluorido [TiF6]2-

chloride Cl- Chlorido [CoCl4]2-

Bromide Br- Bromido [CoBr4]2-

iodide I- Iodido [PtI4]2-

hydroxide OH- Hydroxido [Fe(OH)(H2O)5]2+

methoxide CH3O- Methoxido [Ti(OCH3)4]

cyanide CN- Cyanido [Fe(CN)6]4-

Blackman p 559



EXAMPLES: Look at the Blackboard pages on nomenclature

(i) [Pt(NH3)4]Cl2 tetraammineplatinum(II) chloride

(ii) K4[Fe(CN)6] potassium hexacyanidoferrate(II)

(iii) [Cr(OH2)4Cl2]Cl tetraaquadichloridochromium(III) chloride

(iv) [Ni(en)3]Br2 tris(1,2-diaminoethane)nickel(II) bromide

the prefix is tris (NOT tri); the ligand already has a multiple prefix (di-) in its name

di- (bis-), tri- (tris-), tetra- (tetrakis-), penta- (pentakis-), hexa- (hexakis-)



Name the following complex:- [Ni(NH3)6](SO4)

Hexaamminenickel(II) sulfate (or sulphate)

Convert the following formula to a name:- Na2[NiCl4]

Sodium tetrachloridonickelate(II)



Why are metal complexes coloured????????



Experiment Four: How Sweet Colourless aspartame complexes Cu2+(aq)

to form a blue-coloured complex. The amount of complex formed is measured

using a UV-vis spectrometer, enabling the concentration of aspartame in a sweetener

sachet to be determined.



Crystal Field Model A simple approximation – treats the coordinate bonds as electrostatic in origin. The ligand donor atom (the Lewis base) typically has a high electronegativity (N, O, P, S, halogen). The partial negative charge results in an electrostatic attraction to the positively charged metal ion.

δ-

δ+

δ+

δ+

Blackman p 565-570

N H H

H



Blackman p 566

The five 3d-orbitals

Two point along the axes dz

2 dx2-y

2

Three point between the axes dxy dxz dyz



Blackman p 567



1. In the free ion (no ligands) the five d-orbitals are of the same energy (degenerate). 2. In the presence of the ligands (negative point charges) the d-orbitals split into two sets with different

energies. 3. The degree and nature of the energy splitting depends on the coordination geometry and the type of ligands.

dx2-y

2 dz2

eg set – higher energy

dxy dyz dxz t2g set – lower energy

In an octahedral

complex . . .



d-orbital splitting – Octahedral complexes dz

2 and dx2-y

2 orbitals point towards the ligands (larger repulsion) so the energies are raised considerably dxy, dxz, dyz orbitals point between the ligands, the energies are essentially unchanged

Blackman p 567

∆o 3 t2g orbitals

2 eg orbitals



Tetrahedral complexes

dz2 and dx

2-y

2 orbitals point between the ligands (energies are virtually unaffected) dxy, dxz, dyz orbitals point towards the ligands, their energies are raised. This is the reverse of what occurs for octahedral complexes.

Blackman p 569



http://www.dlt.ncssm.edu/tiger/diagrams/complexions



Tetrahedral complexes

Blackman p 569

Tetrahedral

∆t

3 t2 orbitals

2 e orbitals



t2g

eg

e

t2

Important that you understand and can use these diagrams!!!!

∆t

∆o

Octahedral Tetrahedral



Electronic Configurations

The Pauli exclusion principle is still obeyed (maximum of 2 electrons per orbital, opposite spins). Assuming d-orbital splitting is small (weak field case) the electronic configuration is governed by Hund’s rule (maximum number of unpaired spins), with orbitals filled from lowest to highest energy.

Blackman p 568

The most stable (ground state) d-orbital electronic configuration is a compromise between i. Occupying the lowest energy orbitals AND ii. minimising d electron–d electron

repulsion.



d1


d2



d3




Tetrahedral

d5

d6 d7

d4



d8 d9

d10

Tetrahedral



No ambiguity in the way these (d1, d2, d3) orbitals fill

But for d4 - d7 there are two possibilities, depending on ∆o size.

Blackman p 568 Back to octahedral complexes . . . .



The energy difference between the t2g and eg sets of orbitals is relatively small – weigh the energetic advantage of placing the 4th electron in a low energy t2g orbital against the energetic disadvantage that results from placing an electron in an already occupied orbital. This energy is called the spin pairing energy (P) The magnitude of P compared with Δo determines which electron configuration is favoured. • If P>Δo then the energetically favourable electron configuration is that in which the 4th electron occupies an eg orbital (called a high spin configuration) • If P<Δo the complex adopts a low spin configuration, that is paired with another electron in the t2g orbital.

Blackman p 568



Blackman p 568

high spin low spin



d5

d6

high spin low spin



d7

d8

?

high spin low spin



d8 d9

d10



Spectrochemical series – whether Δ is big or small

The magnitude of the d-orbital splitting (either Δo or Δt) is dependent on the type of ligand The (empirical) ordering of the ligands is governed by the ligand donor atom - It is known as the spectrochemical series

CN- > phen ~ NO2- > en > NH3 ~ py > H2O > C2O4

2- > OH- > F- > S2- > Cl- > Br- > I-

the order of ligand field strength (representing a decrease in ∆o) for common ligands is approximately

big Δo favours low spin

small Δo favours high spin

Blackman p 572



Strong and Weak Field Ligands

Hund’s rule predicts the maximum number of unpaired (parallel) electron spins for a set of degenerate (or similar energy) orbitals Strong field ligands result in a large energy (Δo) difference between the dxy,dxz,dxy and dx

2-y

2 sets Weak field ligands result in a small energy (Δo) difference between the dxy,dxz,dxy and dx

2-y

2 sets

Blackman p 572

(Strong and weak field does NOT refer to the strength of the metal ion-ligand bonds, but to the size of Δo) http://www.naomilkoffman.com/2012/03/25/the-

opposite-party/



Strong field e.g. [Fe(CN)6]3-

Weak field e.g. [Fe(H2O)6]3+


2- > OH- > F- > S2- > Cl- > Br- > I-

Vitally Important!!!!

Blackman p 572

high spin low spin



High and low spin complexes

The most stable (or ground state) configuration is dictated by a compromise between:- • the electron occupancy of lower energy orbitals • the inter-electronic repulsion

Blackman p 568

For d4 – d7 electronic configurations there are two possible electronic ground states with either the . . . . - minimum number of unpaired electrons (low spin) OR - maximum number of unpaired electrons (high spin).



Tetrahedral Complexes Unlike the octahedral complexes, low spin tetrahedral complexes are NOT found. Why? Δt = 4/9Δo (for the same type of ligands) Fewer donor atoms, and poorer overlap with d orbitals Δt is never large enough to enforce electron pairing



Magnetism

Both d6 systems, what are the differences between them?

Blackman p 574



d6 system No unpaired electrons DIAMAGNETIC Remember: It is “dia” to have all paired electrons.

d6 system 4 unpaired electrons PARAMAGNETIC

Important terms!!!!!



Gouy balance schematic. The weight of the sample in a glass tube is measured with or without a magnet present.

Magnetic Susceptibility balance



PARAMAGNETIC DIAMAGNETIC

REPELLED ATTRACTED



MAGNETISM

The number of unpaired electrons (N) - the degree of interaction with magnetic fields We use a term called magnetic susceptibility, given by a numeric quantity called the magnetic moment (μeff) or spin only magnetic moment (μso)

µ eff B MN N= +( ) . .2

B. M. = Bohr magnetons Important formula!!!!!! This is on the data sheet for the exam.

Blackman p 574



( 2) . .eff N N B Mµ = +

N is the number of unpaired electrons Number of unpaired

electrons µeff B.M.

One

Two

Three

Four

Five

2 (2 2) 8 2.83× + = =

1 (1 2) 3 1.73× + = =

3 (3 2) 15 3.87× + = =4 (4 2) 24 4.90× + = =

5 (5 2) 35 5.92× + = =



consider: - octahedral, high spin d4 t2g

3 eg1

N = 4

consider: - octahedral, low spin d4 t2g

4 eg N = 2

µ eff

B M

N N= +

= +=

( )

( ). . .

2

2 2 22 83

µ eff

B M

N N= +

= +=

( )

( ). . .

2

4 4 24 90



Colours of Transition Metal Complexes

Blackman p 570



HEXAAQUA COMPLEXES – [M(H2O)6]2+

Mn(II) Fe(III) Co(II) Ni(II) Cu(II) Zn(II)

Colours of Transition Metal Complexes



Most transition metal complexes are coloured. By definition that means they absorb electromagnetic radiation somewhere in the visible region of the spectrum (400 – 700 nm) Light not absorbed remains to be detected by our eyes.

Blackman p 570





Complementary colours

Light not absorbed remains to be detected by our eyes i.e. the light transmitted (seen by our eyes) is complementary to the light absorbed; opposite on the colour wheel



Important concept!!!!

Sample absorbs all light except that coloured green. Green is perceived.

Sample absorbs violet, red and orange light. Blue, green and yellow light pass through. Green is perceived.



Blackman p 570

This diagram will be provided in the exam. You need to understand how to use it.



Transition metal ions and gem stones

Emeralds are derived from the mineral bery which is beryllium aluminium silicate

3BeO.Al2O3.6SiO2

Some of the Al3+ ions in beryl replaced by Cr3+ and the characteristic colour of

emerald results.

3BeO.(Cr2O3).Al2O3.6SiO2



Gem Stones



[Ti(H2O)6]3+

d1 system



When white light shines on a filter that absorbs in the yellow-green region, the emerging light is violet. Because the complex ion [Ti(H2O)6]

3+ absorbs yellow-green light, a solution of it is violet.

Important concept



Approximate relationship of wavelength of Visible light absorbed to colour observed

Absorbed wavelength in nm (colour)

Observed colour

400 (violet) Greenish yellow 450 (blue) Yellow 490 (blue-green) Red 570 (yellow-green) Violet 580 (yellow) Dark blue 600 (orange) Blue 650 (red) green

Colour(s) absorbed

Absorbing species

Colour(s) transmitted

[CrCl2(H2O)4]+

CrO42-

[Cr(H2O)6]3+

Cr2O72-

[Cu(NH3)4]2+



Several octahedral complexes of Cr3+ and their colours

isomer Colour [Cr(H2O)6]Cl3 Violet

[Cr(H2O)5Cl]Cl2 Blue-green [Cr(H2O)4Cl2]Cl Green [Cr(NH3)6]Cl3 Yellow

[Cr(NH3)5Cl]Cl2 Purple [Cr(NH3)4Cl2]Cl violet

The ligands influence the colours of transition metal complexes. WHY?



Absorption of light must result in some energetic change to the system (excitation). In transition metal complexes, visible light absorption excites electrons from lower energy d orbitals to higher energy d orbitals. Energy of absorbed light (λ) for one photon is inversely proportional to the wavelength of the light.

E hc=λ

Energy of light must match the d-orbital splitting energy for light to be absorbed.

/ Jouleso t

hcλ

∆ =

See page 571 Blackman

h = Planck’s constant c = speed of light in a vacuum



d1

hν

Ground state Excited state

[Ti(H2O)6]3+

Blackman p 570



hν

Blackman p 570



d1

Recall that Δt = 4/9 Δo. So tetrahedral complexes must absorb at longer wavelengths than octahedral analogues (with similar ligand fields) i.e. it takes less energy to promote an electron from a lower energy d-orbital to a higher energy d-orbital.



A complex in an electronically excited state is very short lived (picoseconds pico =10-12) It relaxes to the ground state by releasing energy in the form of heat (molecular vibrations) - usually does not emit light.

Blackman p 570



The magnitude of Δo is dependent on the ligand field strength It increases for strong field ligands. As Δo increases, more energy is required to promote electrons, meaning shorter wavelength.

CN- > H2O > F-

More energy, shorter wavelength Less energy, longer wavelength






2- > OH- > F- > S2- > Cl- > Br- > I-

nm

High Energy; UV light

Lower energy



WAVELENGTH (nm)

COLOUR ABSORBED

COMPLEMENTARY COLOUR

Δo (kJ mol–1)

>720 Infrared Colourless <165 720 Red Green 166 680 Red-orange Blue-green 176 610 Orange Blue 196 580 Yellow Indigo 206 560 Yellow-green Violet 214 530 Green Purple 226 500 Blue-green Red 239 480 Blue Orange 249 430 Indigo Yellow 279 410 Violet Lemon-yellow 292

<400 ultraviolet Colourless >299

∆o t

hc/ =

λJoules (for one photon)

/

hcA1000o t λ

∆ = kJ mol-1 (for one mole of photons)



Consider that an octahedral Cr(III) complex has an absorption band at 486 nm. Calculate the magnitude of the crystal field splitting for this complex.

34 8 -119

9

(6.626 10 J s)(2.998 10 m s ) 4.087 10 J(486 10 m)

hcλ

−−

−

× ×∆ = = = ×

×

These energies are usually expressed in units of kJ mol-1 so we need to convert this

19 23 1 11 kJ4.087 10 J 6.02 10 mol 246.1 kJ mol1000 J

− − −∆ = × × × × =

Blackman p 571

Avogadro constant

Divide the number in Joules by 1000 to convert to kJ.



HEXAAQUA COMPLEXES – [M(H2O)6]2+

Mn(II) Fe(III) Co(II) Ni(II) Cu(II) Zn(II)



Zn(II)

Blackman p 572

Colourless – Why?



Zn(II) is d10

d10 The eg orbitals are already full – there is no space to fit a promoted electron.



Mn(II)

Blackman p 572

Very, very pale pink. Why?



d5

High spin, d5 This excited state can’t exist – can’t have two electrons with all 4 quantum numbers the same – called spin forbidden.

This configuration can’t occur!

X



Many reactions of transition metal complexes are reversible – that is they are equilibrium reactions and the position of the equilibrium can be altered by changes in the concentration of reactants or products, or other factors like temperature.

The Portuguese weather rooster changes colour with the weather. Portions are coated in cobalt chloride, which respond to changes in humidity. In dry weather, it is blue, due to formation of [CoCl4]2-. In humid conditions, pink trans-[CoCl2(H2O)4] is formed.

[CoCl4]2- + 4H2O(l) trans-[CoCl2(H2O)4] + 2Cl-(aq)

Equilibrium p. 561 Blackman



Blackman p 561

Le Chatelier’s Principle: If an outside influence upsets an equilibrium, the system undergoes a change in a direction that counteracts the disturbing influence and, if possible, returns the system to equilibrium. (p. 364 Blackman)

[CoCl4]2- + 4H2O(l) trans-[CoCl2(H2O)4] + 2Cl-(aq)

If water is removed or Cl- ions added, the equilibrium will move to the left, favouring formation of blue [CoCl4]2-.

If water is added or Cl- ions removed, the equilibrium will move to the right, favouring formation of pink trans-[CoCl2(H2O)4]2-.

pink blue



Equilibrium Constant For the general reaction

The equilibrium constant K for the reaction is given by

ba

dc

[B][A][D][C]K =

If K is large, the equilibrium lies to the right (products); If K is small, the equilibrium lies to the left (reactants). Different names, “formation constant, equilibrium constant, association constant, binding constant” all meaning the same thing

Blackman p 547



The “stability of a complex in solution” refers to the degree of association between

the two species involved in the state of equilibrium

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

The magnitude of the stability (or formation constant) for the

association quantitatively expresses the stability

2]3 6

62 362

[Ni(NH )K[Ni(H O) ][NH ]

+

+=

K is an “equilibrium constant”

Blackman p 547



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

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



separate steps - stepwise reactions

[Ni(H2O)6]2+ + NH3 [Ni(OH2)5(NH3)]2+ + H2O [Ni(OH2)5(NH3)]2+ + NH3 [Ni(OH2)4(NH3)2]2+ + H2O

[Ni(OH2)4(NH3)2]2+ + NH3 [Ni(OH2)3(NH3)3]2+ + H2O

Blackman p 561

[Ni(OH2)3(NH3)3]2+ + NH3 [Ni(OH2)2(NH3)4]2+ + H2O

[Ni(OH2)2(NH3)4]2+ + NH3 [Ni(OH2)(NH3)5]2+ + H2O

[Ni(OH2) (NH3)5]2+ + NH3 [Ni(NH3)6]2+ + H2O



overall we can write:-

[Ni(H2O)6]2+ + xNH3 [Ni(OH2)6-x(NH3)x]2+ + H2O

22 3

22 36

6[Ni(H O) (NH ) ][Ni(H O) ][NH ]

xx x

xβ+

+−=

β - Cumulative formation constants

[Ni(H2O)6]2+ + 3NH3 [Ni(OH2)3(NH3)3]2+ + H2O

22 3 33

3 322 36

[Ni(H O) (NH ) ][Ni(H O) ][NH ]

β+

+=

e.g.

Blackman p 561



and, in general

Important concept

Blackman p 561-562

β3 = K1 x K2 x K3

βk = K1 x K2 x K3 x . . . . x Kk

The magnitude of βn for a complex is a measure of how far the formation reaction for that complex proceeds towards completion. If βn is small the complex shows little tendency to form; if βn is large the complex forms almost completely.



Ligand Metal Equilibrium n βn

NH3 Co(II) Co2+(aq) + 6NH3(aq) [Co(NH3)6]2+

(aq) 6 5.0 × 104

Ni(II) Ni2+(aq) + 6NH3(aq) [Ni(NH3)6]2+

(aq)

6 2.0 × 108

en Co(II) Co2+(aq) + 3en(aq) [Co(en)3]2+

(aq)

3 1.0 × 1014

Ni(II) Ni2+(aq) + 3en(aq) [Ni(en)3]2+

(aq)

3 4.1 × 1017

Blackman p. 562

Conclusion 1: Identity of the metal affects stability. Here, Ni2+ complexes of the same ligand are more stable than Co(II) complexes. Cu(II) complexes are generally the most stable of the first row divalent (2+) transition metal ions, when comparing complexes with the same ligands. (CHEM2050 Irving Williams series Mn(II) < Fe(II) < Co(II) < Ni(II) < Cu(II) > Zn(II) )



Conclusion 2: Ligands can show different affinities for various oxidation states of the same metal ion.


NH3 Co(II) Co2+(aq) + 6NH3(aq) [Co(NH3)6]2+

(aq) 6 5.0 × 104

Co(III) Co3+(aq) + 6NH3(aq) [Co(NH3)6]3+

(aq)

6 4.6 × 1033

en Co(II) Co2+(aq) + 3en(aq) [Co(en)3]2+

(aq)

3 1.0 × 1014

Co(III) Co3+(aq) + 3en(aq) [Co(en)3]3+

(aq)

3 5.0 × 1048

Cobalt(III) is favoured in each case.

Blackman p 562



Compare the numbers for complexes of monodentate (NH3), bidentate (en) and hexadentate (EDTA) ligands: Ligand Metal Equilibrium n βn

NH3

Ni(II) Ni2+(aq) + 6NH3(aq) [Ni(NH3)6]2+

(aq)

6 2.0 × 108

en

Ni(II) Ni2+(aq) + 3en(aq) [Ni(en)3]2+

(aq)

3 4.1 × 1017

edta Ni(II) Ni2+

(aq)+EDTA4− (aq) [Ni(EDTA)]2−

(aq) 1 3.6 × 1018

The bidentate ligand en results in a more stable complex (larger βn) than the monodentate ligand NH3. The hexadentate edta ligand results in the most stable complex (largest βn), compared to the en and NH3 complexes. Conclusion 3: The number of chelate rings formed affects the stability of the complex. More chelate rings = more stable.

Blackman p 563



Complexes containing chelate rings are usually more stable than similar

complexes without chelate rings. Generally, as the number of chelate rings increases, so does the stability.

termed - THE CHELATE EFFECT

Important concept

Blackman p 563



[Ni(NH3)6]2+(aq) + H2O(l) [Ni(NH3)5(OH2)]2+ + NH3(aq)

Compare: - A monodentate system

With a chelate

Blackman p 563




EDTA Co(II) Co2+(aq)+EDTA4−

(aq) [Co(EDTA)]2−(aq)

1 2.8 × 1016

EDTA Co(III) Co3+(aq)+EDTA4−

(aq) [Co(EDTA)]−(aq) 1 2.5 × 1041

For EDTA4− we see the effect of both the charge on the metal ion and the chelate effect.



NOTE: The value of βn DOESNT tell us anything about rate of reaction! [Co(NH3)6]3+ + 6H3O+ [Co(OH2)6]3+ + 6NH4

+

The stability constant for this reaction is 1025, so at equilibrium [Co(OH2)6]3+ will be strongly favoured in acidic solution. However, [Co(NH3)6]3+ is stable in 1 M acid for several days. It is kinetically inert to ligand substitution. Cr(III) and low spin Co(III) octahedral complexes are kinetically inert and undergo ligand exchange reactions slowly. (Why inert? Neither have electrons in the eg orbitals, and the t2g orbitals are half or completely filled).

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CHEM1010 Transitio Metal Lec Notes

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