Electronic Configuration and Properties of the Transition Elements

They are often called ‘transition elements’ because their position in the periodic table is between the s – block and p – block elements. Their properties are transitional between the highly reactive metallic elements of the s – block, which typically form ionic compounds, and the elements of the p – block, which are largely covalent. In the s – and p – blocks, electrons are added to the outer shell of the atom. In the d – blocks, electrons are added to the penultimate shell, expanding it from 8 to 18 electrons. Typically, the transition elements configuration and since the d – shell is complete, compounds of these elements are not typical and show some differences from the others. The electrons make up three complete rows of ten elements and an incomplete fourth row. The position of the incomplete fourth series is discussed with the f – block.

METALLIC CHARACTER

In the d – block elements the penultimate shell of electrons is expanding. Thus they have many physical and chemical properties in common. Thus, all the transition elements are metals. They are therefore good conductors of electricity and heat; have a metallic luster and are hard, strong and ductile. They also form alloys with other metals.

VARIABLE OXIDATION STATE

One of the most striking features of the transition elements is that the elements usually exist in several different oxidation states. Furthermore, the oxidation states change in units of one, e.g. Fe³⁺ and Fe²⁺, Cu²⁺ and Cu⁺.

The oxidation states shown by the transition elements may be related to their electronic structures. Calcium, the s – block element preceding the first row of transition elements, has the electronic structure.

Ca 1s²2s²2p⁶3s²3p⁶4s²

It might be expected that the next ten transition elements would have this electronic arrangement with from one to ten d electrons added in a regular way: 3d¹, 3d², 3d³…3d¹⁰. This is true except in the cases of Cr and Cu. In these two cases, one of the s electrons moves into d shell, because of the additional stability when the d orbitals are exactly half filled or completely filled.

Transition elements

Thus, Sc could have an oxidation number of (+11) if both s electrons are used for bonding and (+III) when two s and one d electrons are involved. Ti has an oxidation state (+II) when both s electrons are used for bonding, two d electrons are used. Similarly, V shows oxidation numbers (+II), (+III), (+IV) and (+V). In the case of Cr, by using the single s electron for bonding, we get an oxidation number of (+I): hence by using varying numbers of d electrons oxidation states of (+II), (+III), (+IV), and (+V) and (+VI) are possible. Mn has oxidation states (+II), (+III), (+IV), (+V), (+VI) and (+VII). Among these first five elements, the correlation between electronic structure and minimum and maximum oxidation states in simple compounds is complete. In the highest oxidation states of theses first five elements, all of the s and d electrons are being for bonding. Thus, the properties depend only on the size and valency, and consequently show some similarities with elements of the main groups in similar oxidation states. For example, SO²₄– (Group 16) and CrO²₄– (Group 6) are isostructural, as are SiCl₄ (Group 14) and TiCl₄ (Group 4).

Once the d⁵ configuration is exceeded i.e in the last five elements, the tendency for all the d electrons to participate in bonding decreases. Thus, Fe has a maximum oxidation state of (+VI). However, the second and third elements in this group attain a maximum oxidation state of (+VIII), in RuO₄ and OsO₄. This difference between Fe and the other two elements Ru and Os is attributed to the increased size.

These facts may be conveniently memorized, because the oxidation states form a regular ‘pyramid’ as shown in Table 18.2. Only Sc (+II) and Co(+V) are in doubt. The oxidation number of all elements in the elemental state is zero. In addition, several of the elements have zero-valent and other low-valent states in complexes. Low oxidation states occur particularly with π bonding ligands such as carbon monoxide and dipyridyl.

Similar but not identical pyramids of oxidation states are found on the second and third rows of transition elements. The main differences are as follows:

In Group 8 (the iron group) the second and third row elements show a maximum oxidation state of (+VIII) compared with (+VI) for Fe.

The electronic structures of the atoms in the second and third rows do not always follow the pattern of the first row. The structures of Group 10 elements:

Ni 3d⁸ 4s²

Pd 4d¹⁰ 5s⁰

Pt 5d⁹ 6s¹

Since a full shell of electrons is a stable arrangement, the place where this occurs is of importance.

The d levels are complete at copper, palladium and gold in their respective series.

Ni Cu 3d¹⁰ 4s¹ Zn 3d¹⁰ 4s2

Pd 4d¹⁰ 5s⁰ Ag Cd 3d¹⁰ 4s²

Pt Au 5d¹⁰ 6s¹ Hg 3d¹⁰ 4s²

Even though the ground of the atom has a d¹⁰ configuration, Pd and the coinage metals Cu, Ag and Au behave as typical transition elements. This is because on their most common oxidation states Cu (II) has a d⁹ configuration and Pd (II) and Au (III) have d⁸ configurations, that is they have an incompletely filled d level. However, in zinc, cadmium and mercury, the ions Zn²⁺, Cd²⁺ and Hg²⁺ have d10 configuration. Because of this, these elements do not show the properties characteristics of transition metals.

Stability of the Various Oxidation States

Compounds are regarded as stable if they exist a room temperature, are not oxidized by air, are not hydrolysed by water vapour and do not disproportionate or decompose at normal temperatures. Within each of the transition Groups 3 – 12, there is a difference in stability of the various oxidation states that exist. In general, the second and third row elements exhibit higher coordination numbers, and their higher oxidation states are more stable than the corresponding first row elements. This can be seen more than the corresponding first row elements. This can be seen from Table. This gives the oxides and halides of the first, second and third row transition elements. Stable oxidation states form oxides, fluorides, chlorides, bromides and iodides. Strongly reducing states probably do not form fluorides and/or oxides, but may well form the heavier. Conversely, strongly oxidizing states form oxides and fluorides, but not iodides.

Below are some oxides and halides of the Transition elements

Formation of Complexes By the Transition Elements

The transition elements have an unparalleled tendency to form coordination compounds with Lewis bases; that is with groups which are able to donate an electron pair. These groups are called ligands. A ligand may be a neutral molecule such as NH3, or an ion such as Cl ^– or CN ^–. Cobalt forms more complexes that any other element, and forms more compounds than any other element except carbon.

Co³⁺ + 6NH³ [Co(NH³)₆]³⁺

Fe²⁺ + 6CN ^– [Fe(CN)₆]^{4 –}

In Table, the most stable compounds are bold, unstable compounds are in parenthesis, h indicates hydrated oxides, g indicates that it occurs only as a gas, m indicates metal – metal bonding, c indicates cluster compounds, x indicates mixed oxide and d indicates that it disproportionates.

The ability to form complexes is in marked contrast to the s – and p – block elements which form only a few complexes. The reason transition metals are so good at forming complexes is that they have small, highly charged ions and have vacant low energy orbitals to accept lone pairs of electrons donated by other groups or ligands. Complexes where the metal is in the (+III) oxidation state are generally more stable than those where the metal is in the (+II) state.

Some metal ions form their most stable complexes with ligands in which the donor atoms are N, O or F. Such metal ions include Group 1 and 2 elements, the first half of the transition elements, the lanthanides and actinides, and the p – block elements except for their heaviest member. These metals are called class – a acceptors, and correspond to ‘hard’ acids.. In contrast, the metals Rh, Ir, Pd, Pt, Ag, Au and Hg form their most stable complexes with the heavier elements of Group 15, 16 and 17. These metals are called class – b acceptors, and corresponds to ‘soft acids’ form complex with both types of donors and are thus ‘ intermediate’ in nature, these are shown (a/b) in Table below.

Class – a and Class – b acceptors

(a)

(a/b)

(a)

(a/b)

(b)

(a/b)

(a)

(a/b)

(b)

(a/b)

(a)

SIZE OF ATOMS AND IONS

The covalent radii of the elements decrease from left to right across a row in the transition series, until near the end when the size increases slightly. On passing from left to right, extra protons are placed in the nucleus and extra orbital electrons are added. The orbital electrons shield the nuclear charge incompletely (d electrons shield less efficiently than p – electrons, which in turn shield less effectively than s electrons).

Atoms of the transition elements are smaller than those of the Group 1 or 2 elements in the same horizontal period. This is partly because of the usual contraction in size across a horizontal period discussed above, and partly because the orbital electrons are added to the penultimate d shell rather than to the outer shell of the atom.

The transition elements are divided into vertical groups of three (triads) or sometimes four elements, which have similar electronic structures. On descending one of the main groups of element in the s – and p – blocks, the size of the atoms increases because extra shells of electron are present. The elements in the first group in the d block (Group 3) show the expected increase in size Sc – Y – La. However, in the subsequent Groups (3 – 12), there is an increase in radius of 0.1 – 0.2A between the first and second member, but hardly any increase between the second and third elements. This trend is shown both in the covalent radii and in the ionic radii. Interposed between lanthanium and hafnium are the 14 lanthanide elements, in which the antepenultimate 4f shell of electrons is filled.

Covalent radii of the transition elements (A)

1.57

1.74

1.44

1.32

1.22

1.17

1.16

1.15

1.17

1.25

2.16

1.91

1.62

1.45

1.34

1.29

–

1.24

1.25

1.28

1.34

1.41

2.35

1.98

1.69

1.44

1.34

1.30

1.28

1.26

1.29

1.34

1.44

The effect of the lanthanide contraction or ionic radii

Ca²⁺ 1.00 Sc¹⁺ 0.745 Ti⁴⁺ 0.605 V³⁺ 0.64

Sr²⁺ 1.18 Y³⁺ 0.90 Zr⁴⁺ 0.72 Nb³⁺ 0.72

Ba²⁺ 1.35 La³⁺ 1.032 Hf⁴⁺ 0.71 Ta³⁺ 0.72

There is a gradual decrease in size of the 14 lanthanide elements from cerium to lutetium. This is called the lanthanide contraction. The lanthanide contraction cancels almost exactly covalent radius of Hf and the ionic radius of Hf⁴⁺ are actually smaller than the corresponding values for Zr. The covalent and ionic radii of Nb are the same as the values for Ta. Therefore, the second and third row transition elements have similar radii. As a result, they also have similar lattice energies, salvation energies and ionization energies. Thus, the differences in properties between the first row and second row elements are much greater than the differences between the first row and second row elements. The effects of the lanthanide contraction are less pronounced towards the right of the d block. However, the effect still shows to a lesser degree in the p block elements which follow.

DENSITY

The atomic volumes of the transition elements are low compared with elements in neighboring Group 1 and 2. This is because the increased nuclear charge is poorly screened and so attracts all the electrons more strongly. In addition, the extra electrons added occupy inner orbitals. Consequently, the densities of the transition metals are high. Practically all have a density greater than 5 g cm^-3. (The only exceptions are Sc 3.0g cm^-3 and Y and Ti 4.5g cm^-3). The densities of the second and third row values are even higher; (See Appendix D). The two elements with the highest densities are osmium 22.57g cm^-3 and iridium 22.61g cm^-3. To get some feel for how high this figure really is, a football made of osmium or iridium measuring 30cm in diameter would weigh 320kg or almost one third of a tonne!

MELTING AND BOILING POINTS

The melting and boiling points of the transition elements are generally very high (see Appendices B and C). Transition elements typically melt above 1000^oC. Ten elements melt above 2000^oC and three melt above 3000^oC (Ta 3000^oC, W 3410^oC and Re 3180^oC). There are a few exceptions. The melting points of La and Ag are just under 1000^oC (920^oC and 961^oC respectively). Other notable exceptions are Zn (420^oC), Cd (321^oC) and Hg which is liquid at room temperature and melts at – 38^oC. The last three behave atypically because the d shell is complete, and d electrons do not participate in metallic bonding. The high melting points are in marked contrast to the low melting points for the s block metals Li (181^oC) and Cs (29^oC).

REACTIVITY OF METALS

Many of the metals are sufficiently electropositive to react with mineral acids, liberating H2. A few have low standard electrode potentials and remain unreactive or noble. Noble character is favored by high enthalpies of sublimation, high ionization energies and low enthalpies of solvation. The high melting points indicate high heats of sublimation. The smaller atoms have higher ionization energies, but this is offset by small ions having high salvation energies. This tendency to noble character is most pronounced for the platinum metals (Ru, Rh, Pd, Os, Ir, Pt) and gold.

IONIZATION ENERGIES

The ease with which an electron may be removed from a transition metal atom (that is, its ionization energy) is intermediate between those of the s – and p – blocks. Values for the first ionization energies vary over a wide range from 541kJ mol^-1 for lanthanum to 1007kJ mol^-1 for mercury. These are comparable with the values for lithium and carbon respectively. This would suggest that the transition elements are less electropositive that Groups 1 and 2 and may form either ionic or covalent bonds depending on the conditions. Generally, the lower valent states are ionic and the high valent state covalent. The first row elements have many more ionic compounds than elements in the second and third rows.

COLOUR

Many ionic and covalent compounds of transition elements are coloured. In contrast, compounds of the s – and p – block elements are almost always white. When light passes through a material, it is deprived of those wavelengths that are absorbed. If absorption occurs in the visible region of the spectrum, the transmitted light is coloured with the complementary colour to the colour of the light absorbed. Absorption in the visible and UV regions of the spectrum is caused by changes in electronic energy. Thus the spectra are sometimes called electronic spectra. (These changes are often accompanied by much smaller changes in vibrational and rotational energy). It is always possible to promote an electron from one energy level to another. However, the energy jumps are usually so large that the absorption lies in the UV region. Special circumstances can make it possible to obtain small jumps in electronic energy which appear as absorption in the visible region.

POLARIZATION

NaCl, NaBr and NaI are all ionic are all colourless. AgCl is also colourless; thus the halide ions Cl –, Br – and I –, and the metal ions Na+ and Ag+, are typically colourless. However, AgBr is pale yellow and AgI is yellow. The colour arises because the Ag= ion polarizes the halide ions. This means that it distorts the electron cloud, and implies a greater covalent contribution. The polarization of ions increases with size: thus I is the most polarized, and is the most coloured. For the same reason Ag₂CO₃ and Ag₃PO₄, are yellow, and Ag₂O and Ag₂S are black.

INCOMPLETELY FILLED d or f SHELL

Colour may arise from entirely different cause in ions with incomplete d or f shells. This source of colour is very important in most of the transition metal ions.

In a free isolated gaseous ion, the five d orbitals are degenerate; that is they are identical in energy. In real life situations, the ion will be surrounded by solvent molecules if it is in a solution, by other ligands if it is in a complex, or by other ions if it is in a crystal lattice. The surroundings groups affect the energy of some d orbitals more than others. Thus the d orbitals are no longer degenerate, and at their simplest they form two groups of orbitals of different energy. Thus in transition element ions with a partly filled d shell, it is possible to promote electrons from one d level to another d level of higher energy. This corresponds to a fairly small energy difference, and so light is absorbed in the visible region. The colour of a transition metal complex is dependent on how big the energy difference is between the two d levels. Thus in turn depends on the nature of the ligand, and on the type of complex formed. Thus the octahedral complex and on [Ni(NH₃)₆]²⁺ is blue, [Ni(H₂O)₆]²⁺ is green and [Ni(NO2)6]4 – is brown red. The colour changes with the ligand used. The colour also depends on the number of ligands and the shape of the complex formed.

The source of colour in the lanthanides and the actinides is very similar, arising from f – f transitions. With the lanthanides, the 4f orbitals are deeply embedded inside the atom, and are all shielded by the 5s and 5p electrons. The f electrons are practically unaffected by complex formation: hence the colour remains almost constant for a particular ion regardless of the ligand. The absorption bands are also narrow.

The s – and p – elements do not have a partially filled d shell so there cannot be any d – d transitions. The energy to promote an s or p electron to a higher energy level is much greater and corresponds to ultraviolet light being absorbed. Thus compounds of s – and p – block elements typically are not coloured.Some compounds of the transition metals are white, for example ZnSO₄ and TiO₂. In these compounds, it is not possible to promote electrons with d level. Zn²⁺ has a d¹⁰ configuration and the d level is full. Ti4+ has a d10 configuration and the d level is empty. In the series Sc(+III), Ti(+IV), V(+V), Cr(+VI), and Mn(+VII), these ions may all be considered to have an empty d shell; hence d – d spectra are impossible and these states become increasingly covalent. Rather than form highly charged simple ions, oxoions are formed TiO²⁺, VO , VO , CrO , and MnO . VO is pale yellow, but CrO is strongly yellow coloured , and MnO has an intense purple colour in solution though the solid is almost black. The colour arises by charge transfer. In MnO , an electron is momentarily transferred from O to the metal, thus momentarily changing O^2– to O^– and reducing the oxidation state of the metal from Mn(VII) to Mn(VI). Charge transfer always produces intense colours since the restrictions between atoms.

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