5.firstrowstransitionelements (1)

ZnSc Ti V Cr Mn Fe Ci Ni Cou

Chemistry of the Elements

5 TRANSITION ELEMENTS

FIGURE 5.1 – First Row Transition Elements

Copyright © Pooran Appadu

Ti [Ar]

Sc [Ar]

V [Ar]

Cr [Ar]

Mn [Ar]


`


Fe [Ar]

Co [Ar]

Ni [Ar]

Cu [Ar]

Zn [Ar]

Fe [Ar]

Fe2+ [Ar]

Fe3+ [Ar]

Mn [Ar]

Mn2+ [Ar]

Mn3+ [Ar]


FIGURE 5.2 – Electronic Configuration Of The Elements



FIGURE 5.3 – Electronic Structure Of The Common Ions Of Iron And Manganese

Sc Ti V Cr Mn Fe Co Ni Cu Zn

Atomic No. 3d14s2

3d24s2

3d34s2

3d54s 3d54s2 3d64s2

3d74s2

3d84s2

3d104s

3d104s2

1st I.E / KJ mol-1

630 660 650 650 720 760 760 740 750 910

At. Radius / nm

0.164 0.147 0.135 0.129 0.137 0.126 0.125 0.125 0.128 0.137

M.P / 0C 1540 1680 1900 1890 1240 1540 1500 1450 1080 420B.P / 0C 2730 3260 3400 2480 2100 3000 2900 2730 2600 910

Density g/cm-3 3.00 4.50 6.1 7.2 7.4 7.9 8.9 8.9 8.9 7.1Common Ox

State / Eo+3 +4,

+3+5, +3

+6, +3

+7, +4, +2

+3, +2

+3, +2

+2 +2, +1

+2

M2+/ M -1.20 -0.91 -1.19 -0.44 -0.28 -0.25 -0.34 -0.76M3+/ M -2.1 -1.2 -0.86 -0.74 -0.28 -0.04 +0.40

FIGURE 5.4 – Physical Properties Of The Transition Elements (First Row)



THE TRANSITION ELEMENTS

These are elements that could be considered:

(a) Positioned within the d or f block(b) Forming atleast one ion with a partially filled d – subshell.

COMMON PROPERTIES

They are all hard metals with high melting points and boiling points.

They show more than one oxidation state in their compounds.

They tend to form coloured compounds and ions.

Many show catalytic activity.

They form complex ions in which molecules called ligands for dative bonds with the metal ion.

For the first row the elements range from scandium to zinc. Two elements do not show multiple oxidation state (Sc and Zn) and they form white rather than coloured compounds and they do not act as a catalyst. The reason being:

Scandium atom loses three electrons to give Sc3+ which has the same electronic configuration as Argon.

Zinc atom loses only the outer 4s electrons to give Zn2+ ions with an electronic structure [Ar] 3d10. In neither case is an ion formed with an incomplete d – subshell and so Sc and Zn are not considered true transition elements.

The first row elements have small and relatively similar atomic radius due to the increasing number of electrons being used to fill d – subshell rather than to add new shells. The atomic radius is very little affected until the point is reached where all the d orbital contain one electron when a second electron is pit into a d orbital, the repulsion forces between the two electrons cause the radius of the atom to increase slightly.

For chromium and copper half filled and filled subshells are favourable energetically, so electrons from 4s are transferred to 3d.

TRENDS ACROSS PERIOD OF TRANSITION ELEMENTS

Between sodium to argon electrons are being added to the outershell and the nuclear charge is increasing by addition of protons. The added electrons shield each other only weakly from the extra nuclear charge, so the atomic radii decrease sharply from Na to Ar. At the same time the electronegativities and ionization energies steadily rise.

In moving from Sc to Zn, the nuclear charge is also increasing, but electrons are being added to an inner d sub shell. These inner d electrons shield the outer 4s electrons from the increasing nuclear charge much more effectively than the outer electrons can shield each other. Consequently, the atomic radii decrease much less rapidly.

Similarly, electronegativities and ionization energies increase from Sc to Zn, but only marginally compared with Na to Ar. The increasing electronegativities from Sc to Zn mean that the element becomes slightly less metallic. This is reflected in the increasing positive electrode potential; Eo.



If a salt of vanadium (V), for example, ammonium vanadate, NH4VO3 is dissolved in acid solution and then shaken with a piece of granulated zinc, a striking series of colour changes occur. The vanadium changes from an oxidation state of +5 when it is yellow to oxidation state of +4 when it is blue to oxidation state of +3 when it is green to oxidation state of +2 when it is violet.

VO2+ VO2+ V3+ V2+

COLOUR

When light energy is absorbed by a substance, an electron in the substance is promoted from an orbital of lower energy to one of a higher energy. If the absorbed frequency is in the visible region of the light spectrum, then the materials appeared coloured.

Transition metal ions are often coloured. The five d – orbitals are degenerate, i.e., they all have the same energy level. In a complex ion, the d – orbitals differ slightly in energy as a result overlapping differently with the ligands. Electrons can jump from one d orbital to another if they absorb energy. The colour of the ion is complementary to the colour absorbed. E.g.

Absorbed Reflected

Yellow green VioletYellow orange BlueYellow Blue violetGreen Red VioletBlue Green OrangeGreen Blue Red

The Sc3+ (aq) ion has no d – electrons and is colourless. In the ion Cu+ and Zn2+, with a d10 configuration, no d – d transition is possible, and these ions are colourless. Different ligands affect the energy levels of the d – orbitals.

For example, [Cu(H2O)4]2+ is blue where as [Cu(NH3)4]2+ is very intense deep blue.

Crystal field theory suggests that the five d orbitals be split into two groups, two of the orbitals (dz2 and dx2y2) having a higher energy and three of the orbitals (dxy, dyz and dxz) having a lower energy. The extent of splitting is given the symbol Δo. The two orbitals with the higher energy, e.g. set and the three with the lower energy are the t2g set.



Ligands at the ends of x, y, z axes have electron density that will interact strongly with electrons in the dz2 and dx – y2. These will have higher energies. The lie directly opposite the lobe.


dz2 dx

2 – y2

dxy

dxz

dyz


FIGURE 5.5 – Ligands Attaching To D Orbitals

FIGURE 5.6

Ligand field theory begins by constructing molecular orbitals from the valency d, s and p orbitals of the central metal atom and from the six ligand orbitals that point along the metal ligand bonds in the octahedral complex. The ligand orbital will be so3 orbital containing a lone pair in NH3, or pz orbital from a Cl- ion on the z axis. There are three types of hybridization found in transition metal complexes.


dz2 dz

2 – y2

Δo

dxy dyz dxz

Ener

gy

eg

t2g


Tetrahedral – d3s Square Planar – dsp2 Octahedral – d2sp3

FIGURE 5.7 – Energy Level Filling For A [CoF6]3- ion.


Anti – Bonding Orbitals

Ligand σ (sigma)

Bonding Orbitals

4p

4s

3d

σp*

σs*

σd*

dxynb dxy

nb dxynb

σd

σp

σs

E ne rg y

Co [CoF6]3- 6F-

Δo

Δt

Δ1

Δ1



dxy dyz dxz

dx2- y

2 dz2

Octahedral

d2sp3Δo

dx2- y

2

dxy

dz2

dxy dyz

Square Planar

dsp2

dxy dyz dxz

dx2- y

2 dz2

Δ3

Tetrahedral

NH3

NH3

NH3NH3

NH3

NH3NH3

NH3

NH3

Cl

Cl


FIGURE 5.8 – Octahedral Structure (Ca(NH3)63+

FIGURE 5.9 – Square Planar Structure [Pt(NH3)2Cl2]


Co

NH3

Pt


Cl

Cl ClCl

FIGURE 5.10 – Tetrahedral [FeCl4]-

STABILITY CONSTANT Kstab

Copper (II) sulphate is blue due to the presence of hydrated Cu2+ ions. A standard test for these ions is to add dilute ammonia solution which produce a clear royal blue solution, the tetraamminecopper (II) ion, Cu(NH3)4

2+. The change is an example of equilibrium reaction:

Cu(H2O)42+ (aq) + 4NH3 (aq) Cu(NH3)4

2+ (aq) + 4H2O (l)

Given that the reaction is done in water, whose concentration remains constant, the equilibrium constant is:

[Cu(NH3)42+ (aq)]

K = [Cu(H2O)4

2+] [NH3 (aq)]4

= 1.2 * 1013 dm12 mol-4 at 25oC

If chloride ions are added (e.g. from HCl) to copper II ions in solution, a yellow colour is produced.

Cu(H2O)42+ + 4Cl- (aq) CuCl4

2- (aq) + 4H2O

[CuCl4 (aq)]Kstab =

[Cu(H2O)42+ (aq)] [Cl- (aq)]4

= 4.2 * 105 dm12 mol-4

By comparing the two constants you could tell what the reaction will be if both substances were in the reacting vessel.

Kstab [Cu(NH3)4]2+ = 13.1 (log 1.4 * 1013)

Kstab [CuCl4]2- = 5.60 (log 4.0 * 105)


Fe


So the blue ammine complex will form over the yellow.

For the formation of CuCl42-, the equilibrium is:

[Cu(H2O)6]2+ (aq) + 4Cl- (aq) CuCl42- (aq) + 6H2O (l)

The concentration of water being taken as constant, so the stability constant is written:

[CuCl4 (aq)]Kstab =

[Cu(H2O)42+ (aq)] [Cl- (aq)]4

The larger the value of this constant, the more stable the complex ion. A ligand may be displaced from a metal ion in the presence of another ligand that can form a complex with a larger stability constant. For example, water maybe be displaced as the ligand from [Cu(H2O)6]2+ ion by ammonia forming [Cu(H2O)2(NH3)]2+ ion. This in turn may be displaced by the polydentate ligand edta4- to form [Cuedta]2-. Finally, edta2- may be displaced by cyanide form of [Cu(CN)4]2-. This series of displacement is as follows:

[Cu(H2O)6]2- [Cu(H2O)2(NH3)4]2+ Cu(edta)2- Cu(CN)42-

Ligand

EquilibriumMm+ + XLL MLx

(m-xL)+Stability Constant

@ 25oC / mol dm-3

NH3 Ag+ + 2NH3 [Ag(NH3)2]+

Cu2+ + 4NH3 [Cu(NH3)4]2+

1.7 * 107

1.4 * 1013

Cl-CO3+ + 6NH3

[Co(NH3)6]3+

Cu2+ + 6Cl- CuCl42-

Hg2+ + 4Cl- HgCl42-

4.5 * 1033

4.0 * 105

1.7 * 1016

CN-Cd2+ + 4CN-

[Cd(CN4)]2-

Fe2+ + 6CN- [Fe(CN4)]2-

Hg2+ + 4CN- [Hg(CN)4]2

7.1 * 1016

1.0 * 1024

2.5 * 1041

TABLE 5.11 – Reactions Of Common Ligands


XSNH3

XSEdta

XSCyanide

: :

N N N

FeFe


The large stability constants for complexes with cyanide ions (CN-) as ligands show why cyanide is such as powerful poison. It binds irreversibly to Fe2+ ions are essential for the transport of oxygen in the blood. Cyanide binds more strongly than oxygen to the Fe2+ ion in haemoglobin and this stops the carriage of oxygen in the blood.

The same explanation goes for carbon monoxide which is a poisonous gas and has a high stability constant.

The electronic structure of carbon monoxide is represented:

C OThe lone pair of electrons on the carbon atom enables it to act as a ligand in the formation of carbonyl complexes e.g. [Cr(CO6)], -[FeCO], N: (CO)4 etc.

Carbon monoxide is a neutral molecule and therefore has no charge. The metal, it combines with has oxidation state of 2+.

O2

H2O


NN N N

N

N

NH

Fe


FIGURE 5.12 – Reactions Of Common Ligands

The mode of oxygen carriage by haemoglobin with a water molecule which is reversibly replaced by an oxygen molecule.


:

N N

C ≡ O Lone pair from carbon attach to Fe2+ on RBC

N


FIGURE 5.13 – Haemoglobin Binding With Carbon Monoxide

LIGAND

NAME

: No2- Nitro:

OCO22-

Carbonate

: ONO-

Nitrate

: CN- Cyano: SCN- Thiocyana

to: NCS- Isothiocya

nato: OH- Hydroxo:OH2 Aqua: NH3 Amine: CO Carbonyl

: NO2- Nitrosyl

TABLE 5.14 – Common Ligands


N NN


Carbon monoxide replaces both H2O and O2 and form a stable bond (ligand).

FIGURE 5.15 – Formation Of A Stable Bond With Carbon Monoxide

VARIABLE OXIDATION STATES

d – subshells are very close in energy level to the 4s subshell, and so it is relatively easy lose electrons. This means that several different ions of the same element are possible by losing different numbers of electrons – all of these different ions are approximately equally stable. The inter conversions between one oxidation state and another are an important aspect of transition metal chemistry, and the interconversions are frequently refuted by colour changes in solutions of the complexions.

Some general observations are the variable oxidation states of the transition metals titanium to copper are given below:

1. +1, +2 or +3 are among the most common oxidation states for each element. +3 is most common from Ti to Cr, then from Mn onwards +2 is more common.

2. Transition metals usually show their highest oxidation state when combined with O2 or F2 or Cl2

the most electronegative elements.

3. The highest oxidation states of all the elements upto and including manganese compounds correspond to the loss in bonding of all the electrons outside the argon core. Thus, the highest oxidation state of titanium is +4 whilst that of manganese is +7. Beyond manganese the d – electrons are held more strongly as a result of increasing nuclear charge, and so the common oxidation states involves the 4s shell only.

4. When the elements exhibit high oxidation state (+4 and above) they do not form simple ions. They either are involved in covalent bonding e.g. (TiO2, TiO4, CrO3, Mn2O7) or they form large ions such as CrO4

2- or MnO4-.


M C ≡ O: :


Element

Common Oxidation States

Equation Of Compounds

Ti +2, +3, +4 Ti2O3, TiO2, TiCl4V +1, +2, +3, +4, +5 V2O3, VCl3, V2O5Cr +1, +2, +3, +4, +5,

+6Cr2O3, CrCl3, CrO3

Mn +1, +2, +3, +4, +5, +6, +7

MnO, MnCl2, MnO2, Mn2O7

Fe +1, +2, +3, +4, +5 FeO, FeCl2, Fe2O3, FeCl3 Co +1, +2, +3, +4, +5 CoO, Co2O, CoCl3Ni +1, +2, +3, +4 NiO, NiCl2Cu +1, +2, +3 Cu2O, CuCl2, CuO, CuCl

TABLE 5.16 – Common Oxidation States Of The Transition Elements (First Row)

Note that according to Aufban principle orbitals are required to be filled in increasing order of energy. In the 4th period it is an experimental fact that the 4s orbital actually has a lower energy than the 3d resulting in configuration. For K = [Ar]4s1 and Ca = [Ar] 4s2. The energy of the 3d orbitals does decline relative to the 4s with increasing nuclear charge. So from Sc the next orbital that is filled is the 3d and lies below that of 4s in energies. Electrons are first removed from the 4s to form ions.

Example: Sc+ = [Ar] 3d1 4s1

3d FillingSc Ti V Cr

3d54s1Mn Fe Co Ni Cu

3d104s1Zn

4d FillingY Zr Nb

4d45s1Mo

4d55s1Tc Ru

4d75s1Rh

4d85s1Pb

4d10Ag

4d105s1Cd

5d FillingLu Hf Ta W Re OS Ir Pt

5d96s1Au

5d106s1Hg

6d FillingLw

TABLE 5.17 – Anomalous Eletronic Configuration Of The Transition Elements

Note the group with Cu, Ag, Au and Cr and Mo, so as to account for complex ions.


5.firstrowstransitionelements (1)

Documents