Transition Elements&Catalysts http://www.chem.ox.ac.uk/vrchemistry/complex/allbottlesmsiedefault.html http://www.gcsescience.com/pt21.htm

TRANSITION METALS High densities,melting and boiling points Ability to exist in a variety of oxidation states Formation of colored ions Ability to form complex ions Ability to act as catalysts Zn and Sc do not share these properties.

THE FIRST ROW TRANSITION ELEMENTS Definition D-block elements forming one or more stable ions with partially filled (incomplete) d-sub shells. The first row runs from scandium to zinc filling the 3d orbitals. Properties arise from an incomplete d sub-shell in atoms or ions

THE FIRST ROW TRANSITION ELEMENTS Metallic properties all the transition elements are metals strong metallic bonds due to small ionic size and close packing higher melting, boiling points and densities than s-block metals K Ca Sc Ti V Cr Mn Fe Co m. pt / °C 63 850 1400 1677 1917 1903 1244 1539 1495

Sc and Zn A transition metal is one which forms one or more stable ions which have incompletely filled d orbitals.

Scandium Scandium has the electronic structure [Ar] 3d14s2. When it forms ions, it always loses the 3 outer electrons and ends up with an argon structure. The Sc3+ ion has no d electrons and so doesn't meet the definition.

Zinc Zinc has the electronic structure [Ar] 3d104s2. When it forms ions, it always loses the two 4s electrons to give a 2+ ion with the electronic structure [Ar] 3d10. The zinc ion has full d levels and doesn't meet the definition either.

Copper By contrast, copper, [Ar] 3d104s1, forms two ions. In the Cu+ ion the electronic structure is [Ar] 3d10. However, the more common Cu2+ ion has the structure [Ar] 3d9. Copper is definitely a transition metal because the Cu2+ ion has an incomplete d level.

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS POTASSIUM 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s1 ‘Aufbau’ Principle In numerical terms one would expect the 3d orbitals to be filled next. However, because the principal energy levels get closer together as you go further from the nucleus coupled with the splitting into sub energy levels, the 4s orbital is of a LOWER ENERGY than the 3d orbitals so gets filled first.

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS CALCIUM 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 As expected, the next electron in pairs up to complete a filled 4s orbital. This explanation, using sub levels fits in with the position of potassium and calcium in the Periodic Table. All elements with an -s1 electronic configuration are in Group I and all with an -s2 configuration are in Group II.

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS SCANDIUM 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d1 With the lower energy 4s orbital filled, the next electrons can now fill p the 3d orbitals. There are five d orbitals. They are filled according to Hund’s Rule. BUT WATCH OUT FOR TWO SPECIAL CASES.

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS TITANIUM 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d2 The 3d orbitals are filled according to Hund’s rule so the next electron doesn’t pair up but goes into an empty orbital in the same sub level. HUND’S RULE OF MAXIMUM MULTIPLICITY

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS VANADIUM 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d3 The 3d orbitals are filled according to Hund’s rule so the next electron doesn’t pair up but goes into an empty orbital in the same sub level. HUND’S RULE OF MAXIMUM MULTIPLICITY

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS CHROMIUM 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s1 3d5 One would expect the configuration of chromium atoms to end in 4s2 3d4. To achieve a more stable arrangement of lower energy, one of the 4s electrons is promoted into the 3d to give six unpaired electrons with lower repulsion.

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS MANGANESE 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d5 The new electron goes into the 4s to restore its filled state.

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS IRON 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Orbitals are filled according to Hund’s Rule. They continue to pair up. HUND’S RULE OF MAXIMUM MULTIPLICITY

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS COBALT 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d7 Orbitals are filled according to Hund’s Rule. They continue to pair up. HUND’S RULE OF MAXIMUM MULTIPLICITY

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS NICKEL 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d8 Orbitals are filled according to Hund’s Rule. They continue to pair up. HUND’S RULE OF MAXIMUM MULTIPLICITY

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS COPPER 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s1 3d10 One would expect the configuration of copper atoms to end in 4s2 3d9. To achieve a more stable arrangement of lower energy, one of the 4s electrons is promoted into the 3d. HUND’S RULE OF MAXIMUM MULTIPLICITY

ELECTRONIC CONFIGURATIONS OF THE FIRST ROW TRANSITION METALS ZINC 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s2 3d10 The electron goes into the 4s to restore its filled state and complete the 3d and 4s orbital filling.

ELECTRONIC CONFIGURATIONS K 1s2 2s2 2p6 3s2 3p6 4s1 Ca 1s2 2s2 2p6 3s2 3p6 4s2 Sc 1s2 2s2 2p6 3s2 3p6 4s2 3d1 Ti 1s2 2s2 2p6 3s2 3p6 4s2 3d2 V 1s2 2s2 2p6 3s2 3p6 4s2 3d3 Cr 1s2 2s2 2p6 3s2 3p6 4s1 3d5 Mn 1s2 2s2 2p6 3s2 3p6 4s2 3d5 Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Co 1s2 2s2 2p6 3s2 3p6 4s2 3d7 Ni 1s2 2s2 2p6 3s2 3p6 4s2 3d8 Cu 1s2 2s2 2p6 3s2 3p6 4s1 3d10 Zn 1s2 2s2 2p6 3s2 3p6 4s2 3d10

To write the electronic structure for Fe3+: Fe 1s22s22p63s23p63d64s2 Fe3+1s22s22p63s23p63d5 The 4s electrons are lost first followed by one of the 3d electrons. The rule is quite simple. Take the 4s electrons off first, and then as many 3d electrons as necessary to produce the correct positive charge.

When d-block elements form ions, the 4s electrons are lost first. For Cr [Ar]4s13d5 is the configuration for chromium atom. It has this rather than two electrons in the 4s as it is more stable like this. When it forms 3+ it has the configuration [Ar]4s03d3 as the 4s will be lost first as they are at a lower energy level. http://www.chemguide.co.uk/inorganic/transition/features.html

VARIABLE OXIDATION STATES Arises from the similar energies required for removal of 4s and 3d electrons maximum rises across row to manganese maximum falls as the energy required to remove more electrons becomes very high all (except scandium) have an M2+ ion stability of +2 state increases across the row due to increase in the 3rd Ionisation Energy THE MOST IMPORTANT STATES ARE IN RED Ti Sc V Cr Mn Fe Co Ni Cu Zn +1 +2 +3 +4 +5 +6 +7 When electrons are removed they come from the 4s orbitals first Cu 1s2 2s2 2p6 3s2 3p6 3d10 4s1 Ti 1s2 2s2 2p6 3s2 3p6 3d2 4s2 Cu+ 1s2 2s2 2p6 3s2 3p6 3d10 Ti2+ 1s2 2s2 2p6 3s2 3p6 3d2 Cu2+ 1s2 2s2 2p6 3s2 3p6 3d9 Ti3+ 1s2 2s 2p6 3s2 3p6 3d1 Ti4+ 1s2 2s2 2p6 3s2 3p6

COLOURED IONS A characteristic of transition metals is their ability to form coloured compounds Theory ions with a d10 (full) or d0 (empty) configuration are colourless ions with partially filled d-orbitals tend to be coloured it is caused by the ease of transition of electrons between energy levels energy is absorbed when an electron is promoted to a higher level the frequency of light is proportional to the energy difference.

The observed colour of a solution depends on the wavelengths absorbed Copper sulphate solution appears blue because the energy absorbed corresponds to red and yellow wavelengths. Wavelengths corresponding to blue light aren’t absorbed. WHITE LIGHT GOES IN SOLUTION APPEARS BLUE ENERGY CORRESPONDING TO THESE COLOURS IS ABSORBED Absorbed colour nm Observed colour nm VIOLET 400 GREEN-YELLOW 560 BLUE 450 YELLOW 600 BLUE-GREEN 490 RED 620 YELLOW-GREEN 570 VIOLET 410 YELLOW 580 DARK BLUE 430 ORANGE 600 BLUE 450 RED 650 GREEN 520

a solution of copper(II)sulphate is blue because red and yellow wavelengths are absorbed blue and green not absorbed white light

Coordination Compound Open Power Point Coordination Chemistry

A Coordination Compound Small size d block ions tend to attract species rich in electrons forming complex ions. Typically consists of a complex ion and counterions (anions or cations as needed to produce a neutral compound): [Co(NH3)5Cl]Cl2 K3Fe(CN)6

Ligands Neutral molecule or ion having a lone electron pair that can be used to form a bond to a metal ion. Monodentate ligand – one bond to a metal ion Bidentate ligand – two bonds to a metal ion Polydentate ligand – more than two bonds to a metal ion

COMPLEX IONS - LIGANDS Formation ligands form co-ordinate bonds to a central transition metal ion Ligands atoms, or ions, which possess lone pairs of electrons form co-ordinate bonds to the central ion donate a lone pair into vacant orbitals on the central species Ligand Formula Name of ligand chloride Cl¯ chloro cyanide NC¯ cyano hydroxide HO¯ hydroxo oxide O2- oxo water H2O aqua ammonia NH3 ammine some ligands attach themselves using two or more lone pairs classified by the number of lone pairs they use multidentate and bidentate ligands lead to more stable complexes

some ligands attach themselves using two or more lone pairs classified by the number of lone pairs they use multidentate and bidentate ligands lead to more stable complexes Unidentate form one co-ordinate bond Cl¯, OH¯, CN¯, NH3, and H2O Bidentate form two co-ordinate bonds H2NCH2CH2NH2 , C2O42-

Coordination Number Number of bonds formed between the metal ion and the ligands in the complex ion. 6 (octahedral shape)and 4 (tetrahedral shape) 2 and 8 (least common, linear)

Coordination Number: 6, 4, 2

when transition metals form coordination complexes, the d-orbitals of the metal interact with the electron cloud of the ligands in such a manner that the d-orbitals become non-degenerate (not all having the same energy.) The way in which the orbitals are split into different energy levels is dependent on the geometry of the complex.

http://www.chemguide.co.uk/inorganic/complexions/colour.html When the ligands bond with the transition metal ion, there is repulsion between the electrons in the ligands and the electrons in the d orbitals of the metal ion. That raises the energy of the d orbitals. However, because of the way the d orbitals are arranged in space, it doesn't raise all their energies by the same amount. Instead, it splits them into two groups.

INCREASING ENERGY / DISTANCE FROM NUCLEUS COPPER 4 4p 4d 4f INCREASING ENERGY / DISTANCE FROM NUCLEUS 3d 4s 3 3p 1s2 2s2 2p6 3s2 3p6 4s1 3d10 One would expect the configuration of copper atoms to end in 4s2 3d9. To achieve a more stable arrangement of lower energy, one of the 4s electrons is promoted into the 3d. HUND’S RULE OF MAXIMUM MULTIPLICITY

http://www.chemguide.co.uk/inorganic/complexions/colour.html The diagram shows the arrangement of the d electrons in a Cu2+ =[Ar]3d9 before and after six water molecules bond with it.

Whenever 6 ligands are arranged around a transition metal ion, the d orbitals are always split into 2 groups in this way : 2 with a higher energy than the other 3. The size of the energy gap between them (shown by the blue arrows on the diagram) varies with the nature of the transition metal ion, its oxidation state (whether it is 3+ or 2+, for example), and the nature of the ligands.

When white light is passed through a solution of this ion, some of the energy in the light is used to promote an electron from the lower set of orbitals into a space in the upper set. http://www.chemguide.co.uk/inorganic/complexions/colour.html

SPLITTING OF 3d ORBITALS Placing ligands around a central ion causes the energies of the d orbitals to change Some of the d orbitals gain energy and some lose energy In an octahedral complex, two go higher and three go lower In a tetrahedral complex, three go higher and two go lower Degree of splitting depends on the CENTRAL ION and the LIGAND The energy difference between the levels affects how much energy is absorbed when an electron is promoted. The amount of energy governs the colour of light absorbed. OCTAHEDRAL TETRAHEDRAL 3d 3d The light that is absorbed corresponds to the energy required to promote a d electron from the lower split level to the higher split level

The magnitude of the splitting of the d-orbitals in a transition metal complex depends on three things: the geometry(shape) of the complex the oxidation state of the metal the nature of the ligands

Complex ion - Iron Water is a common ligand as in hexaaquairon (III) ion, [Fe(H2O)6 3+

COMPLEX IONS - LIGANDS some ligands attach themselves using two or more lone pairs classified by the number of lone pairs they use multidentate and bidentate ligands lead to more stable complexes Multidentate form several co-ordinate bonds HAEM A complex containing iron(II) which is responsible for the red colour in blood and for the transport of oxygen by red blood cells. Co-ordination of CO molecules interferes with the process

COMPLEX IONS - LIGANDS some ligands attach themselves using two or more lone pairs classified by the number of lone pairs they use multidentate and bidentate ligands lead to more stable complexes Multidentate form several co-ordinate bonds

CO-ORDINATION NUMBER & SHAPE the shape of a complex is governed by the number of ligands around the central ion the co-ordination number gives the number of ligands around the central ion a change of ligand can affect the co-ordination number Co-ordination No. Shape Example(s) 6 Octahedral [Cu(H2O)6]2+ 4 Tetrahedral [CuCl4]2- Square planar Pt(NH3)2Cl2 2 Linear [Ag(NH3)2]+

Ligands can be replaced by other ligands. Copper and Ammonia Ligands can be replaced by other ligands. The addition of ammonia to an aqueous solution of copper sulfate gives a deep blue color of tetramine copperII ion. Cl- NH3 [CuCl4]2-  [Cu(H2O)4 ]2+  [Cu(NH3)4 ]2+ H2O H2O

ISOMERISATION IN COMPLEXES GEOMETRICAL (CIS-TRANS) ISOMERISM Square planar complexes of the form [MA2B2]n+ exist in two forms trans platin cis platin An important anti-cancer drug. It is a square planar, 4 co-ordinate complex of platinum.

13.2.7. Catalytic Action of Transition Elements http://www.youtube.com/watch?v=NIZVXVwvsrA Catalysts increase the rate of a chemical reaction without themselves being chemically changed(lower EA). They can be heterogeneous( catalyst is in a different phase from the reactants) or homogeneous( same phase)

Transition metals and their compounds show great catalytic activity It is due to partly filled d-orbitals which can be used to form bonds with adsorbed reactants which helps reactions take place more easily Transition metals can both lend electrons to and take electrons from other molecules. By giving and taking electrons so easily, transition metal catalysts speed up reactions Examples of catalysts IRON Manufacture of ammonia - Haber Process NICKEL Hydrogenation reactions - margarine manufacture MANGANESE IV OXIDE Hydrogen peroxide VANADIUM(V) OXIDE Manufacture of sulphuric acid - Contact Process http://www.gcsescience.com/pt21.htm

Read CC page 59,60. Outline due Monday.

Co in Vitamin B12 Vitamin B12, also called cobalamin, is a water-soluble vitamin with a key role in the normal functioning of the brain and nervous system, and for the formation of blood. It is one of the eight B vitamins. It is normally involved in the metabolism of every cell of the human body, especially affecting DNA synthesis and regulation, but also fatty acid synthesis and energy production. Cobalt is a biological catalyst.

Contact Process Vanadium pentoxide is used in different, industrial processes as catalyst: In the contact process it serves for the oxidation of SO2 to SO3 with oxygen at 440°C. Sulfur dioxide and oxygen then react as follows: 2 SO2(g) + O2(g) ⇌ 2 SO3(g) : ΔH = −197 kJ mol−1 To increase the reaction rate, high temperatures (450 °C), medium pressures (1-2 atm), and vanadium(V) oxide (V2O5) are used to ensure a 96% conversion. Platinum would be a more effective catalyst, but it is very costly and easily poisoned.[ The catalyst only serves to increase the rate of reaction as it has no effect on how much SO3 is produced. The mechanism for the action of the catalyst is: http://www.chemguide.co.uk/physical/equilibria/contact.html

Haber Process N2(g) + 3H2(g)  2NH3 The Haber Process combines nitrogen and hydrogen into ammonia. The nitrogen comes from the air and the hydrogen is obtained mainly from natural gas (methane). Iron is used as a catalyst. N2(g) + 3H2(g)  2NH3

Iron in hemoglobin(iron is a biological catalyst) A hemoglobin molecule is made up of a heme unit covalently bonded to a protein chain. Hemes are most commonly recognized in their presence as components of hemoglobin, the red pigment in blood,

Decomposition of Hydrogen Peroxide – MnO2 Hydrogen peroxide decomposes to oxygen and water when a small amounts of manganese dioxide is added.. 2 H2O2 → 2 H2O + O2

Nitrogen Ethene reacts with hydrogen in the presence of a finely divided nickel catalyst at a temperature of about 150°C. Ethane is produced.

Catalytic Converter Catalytic converters change poisonous molecules like carbon monoxide and various nitrogen oxides in car exhausts into more harmless molecules like carbon dioxide and nitrogen. They use expensive metals like platinum, palladium and rhodium as the heterogeneous catalyst. The metals are deposited as thin layers onto a ceramic honeycomb. This maximises the surface area and keeps the amount of metal used to a minimum. http://auto.howstuffworks.com/catalytic-converter2.htm

Catalytic Converters CO + Unburned Hydrocarbons + O2 CO2 + H2O 2NO + 2NO2 2N2 + 3O2 13.6

13.2.8. Economic Significance of catalysts in Contact and Haber processes http://www.youtube.com/watch?v=5b8dVTucMlY The catalytic properties of these metals are due to their ability to exist in a number of stable oxidation states and the presence of empty orbitals for temporary bond formation. The catalyst is usually a powdered solid and the reactants a mixture of gases. Cc page 66

http://hferrier.co.uk/higher/unit1a/unit1a.htm They increase the rates of reaction and decrease the time needed for a reaction to reach equilibrium. They increase the efficiency of industrial processes and help reduce costs and so increase profits.

So, a catalyst affects the transition state and activation path So, a catalyst affects the transition state and activation path. How does it do this? Typically, by complexing one of the reagents. Complexation by transition metals affords access to a wide variety of oxidation states for the metal. This has the property of providing electrons or withdrawing electrons from the transition state of the reaction. That is, if the transition state is electron rich, then the transition metal might hold some of that electron density and those prevent too much from building up on the reagent. This would then facilitate the reaction. Or the transition metal might undergo formal oxidation/reduction to achieve electron transfer to a substrate, thereby allowing a reaction to occur. This is "complexation and electron storage" taken to the extreme but is a common mechanism in organometallic chemistry. Indeed, a variety of catalytic pathways rely on a two electron transfer between the metal and the substrate (e.g. hydroformylation). It is the ability of the transition metal to be in a variety of oxidation states, to undergo facile transitions between these oxidation states, to coordinate to a substrate, and to be a good source/sink for electrons that makes transition metals such good catalysts.