As we are familiar that elements are categorized on the basis of their electronic configuration into s-block, p-block, d-block and f-block elements. The s and p block elements altogether represent one of the main groups of the elements and are known as main group elements. The d block and the f-block elements symbolize the transition and the inner-transition elements correspondingly. The name transition is given to the elements on the base of their position or place in the periodic table and their properties, that is, they occupy a position between the highly electropositive elements on the left and the electronegative elements on the right. Their properties are as well midway of the s and p block elements.
Here we would like to draw your notice to the fact that a few chemists consider as transition elements only those which, either as neutral atoms or in any of their common oxidation states, have partially filled d-orbitals. On the basis of this definition, the elements Zn, Cd and Hg are excluded from the list of transition elements.
We are familiar that the electronic configuration of the argon atom is ls22s22p63p6. In the atoms of successive elements from potassium to zinc, electrons can either enter 3d or 4s levels. In potassium and calcium atoms the differentiating electrons enter the 4 level; therefore the electronic configuration of calcium can be represented as [Ar]4s2. At scandium the 3d level starts to fill rather than 4p and the resultant configuration of the atoms of transition elements is illustrated in the table shown below. The electronic configuration of the ions can be achieved by eliminating first the outer s electrons of the atom and then the outer d electrons till the total number of electrons removed is equivalent to the charge on the ion. For illustration: The configuration of Ni is [Ar]3d8.
Table: Electronic configurations of free atoms and dipositive ions of the first transition series
Now this might ask as to why the electrons in potassium enter 4s level instead of 3d and then later (from Sc to Cu) why 3d levels are filled before to 4p level.
The radial dependence of the d-orbitals is mainly responsible for this order of filling of electrons in such elements. The below figure exhibits the plot of radial probability functions for a 3d and 4s electron in the hydrogen atom. Let us suppose that the radial probability functions for 3d and 4s electron in a multi electron atom follow the similar pattern as in the hydrogen atom.
Fig: Radial probability functions for 3d and 4s orbitals in hydrogen atom.
Here, ao is the radius of first Bohr orbit is 53.9 pm.
We can observe from the figure above that significant humps in the 4s probability function take place close to the origin, and well inside the maximum of the 3d probability function. This proposes that the 4s electron penetrates considerably into the argon core and spends an appreciable part of its time close to the nucleus. The average nuclear charge experienced via the 4s electron is, thus, higher than that experienced via the 3d electron and therefore after argon, in potassium and calcium the electrons enter the 4s orbital instead of the 3d. As such two electrons are added the nuclear charge is as well increased by two units. Since the 3d orbitals penetrate the 4s orbital more than the 4p orbitals can penetrate the 4s orbital, the total result is that the effective nuclear charge for the 3d orbitals rises abruptly and they now drop well beneath the 4p orbitals to around the level of 4s orbital. Furthermore, as the atomic number increases, the 3d probability maximum gradually moves closer to the core and they continue to drop in energy. The next electron, thus, enters the 3d orbital prior to the 4p orbital. The variation of the energies of the orbitals having increasing atomic number is illustrated very clearly in the figure shown below.
Fig: variation of energy of atomic orbitals-increasing atomic number
This process continues till the whole 3d shell is filled. Therefore at Zn we comprise the configuration [Ar]4s23d10. After that, the subsequent lowest available orbitals are 4p which get filled in the subsequent six elements. This similar sequence of events for the filling of 5s and 4d orbitals is repeated again in the elements following the krypton in second transition series. This series begins with Y and is completed at Cd having the configuration [Kr]4d105s2. After xenon, [Kr]4d105s25p6, the subsequent available orbitals are 4f, 5d, 6s and 6p orbitals. The 4f orbitals are so slightly penetrating with respect to the xenon core that they have barely gained any stability, whereas the more penetrating 6s and 6p levels have gained an excellent deal of stability. Therefore, in the next two elements, electrons are added to 6s orbitals giving Cs and Ba, correspondingly. Though, the 6s electrons don't shield the 4f orbitals efficiently, thus the latter abruptly feel an increase in efficient nuclear charge and therefore suffer a steep drop in energy. At the similar time, by the addition of electrons in the 6s orbital, the 5d orbitals as well drop in energy in the similar manner as the 3d ones. This makes a situation in which 5d and 4f orbitals are of almost similar energy. The subsequent electron in lanthanum therefore enters the 5d orbital, however in the following element cerium; the electronic configuration is [Xe]6s25d14f1.
The electrons then carry on to be added to the 4f orbital till we arrive at ytterbium which has the configuration [Xe]6s24f14. Now by the 6s and 4f shells full, the next lowest levels are the 5d's. Therefore from lutecium onwards, the electrons enter the 5d orbital. This carries on till we reach mercury which consists of the configuration [Xe]6s24f145d10. The electronic configurations of transition elements of 4d and 5d transition series are illustrated in the table shown below.
Table: Electronic configuration of 4d and 5d transition elements
If the filling of orbitals in transition elements occurs via the above scheme, then you might wonder why in the case of some elements example: Cr & Cu (that is, belonging to the first transition sends) and Mo and Ag (that is, belonging to the second transition series) the electronic configuration is represented as [Ar]3d54sl and [Ar]3d104s1 and [Kr]4d105s1, correspondingly. This is due to these configurations are considered to provide more stability to the elements, instead of [Ar]3d44s2 and [Ar] 3d94s2 and [Kr] 4d95s2 correspondingly. This obvious stability can be related by the high stability of exactly half filled and completely filled orbitals. Half-filled and completely-filled orbitals include exchange energy considerably more than the exchange energies related by any other configuration. This exchange energy is the driving force for these configurations to take an electron out of turn in order to accomplish or maintain the half-filled or completely-filled configuration. As well these configurations give the most symmetrical distribution of electrons that suffer the minimum mutual repulsion.
The exchange energy for any configuration is proportional to the total number of possible pairs of electrons having parallel spin in any orbital, that is, Eex = K x P, here 'K' is a constant and 'P' is the number of possible pairs of electrons with parallel spin. If 'n' is the number of electrons having parallel spin for any configuration, 'P' will be equivalent to nC2. Accordingly the values of 'P' for different values of 'n' are represented below:
n 1 2 3 4 5 6 7
p 0 1 3 6 10 15 21
Now, compare the exchange energy for two possible configurations 3d44s2 and 3d54s1 for chromium.
Fig: Exchange energy for two possible configurations
The electrons present in 4s orbital in two configurations contribute nothing to exchange energy as they don't comprise any pair with the parallel spin. Four unpaired d-elections in first configuration can make six pairs of electrons having parallel spin and therefore contribute 6K towards the exchange energy while five unpaired d-electrons in second configuration contribute 10K towards exchange energy as they can comprise 10 combinations of pairs of electrons having parallel spin. This gain of 4K in exchange energy would favor the 3d54s1 configuration for chromium. However you must keep in mind that in accomplishing this configuration, there would be loss of energy in promoting an electron from 4s to 3d orbital. In case of chromium, the gain in exchange energy is greater than the loss in energy and thus, 3d54s1 is the favored configuration.
Likewise you can compare the exchange energies for two possible configurations 3d94s2 and 3d104s1 for copper.
Fig: Exchange energies for copper
The former configuration consists of two sets of electrons having parallel spin - one set consists of five electrons symbolized by upward arrows and the other consists of four electrons symbolized by downward arrows.
These two sets of electrons will contribute 10K and 6K that is, a total 16K towards exchange energy. On the other hand, the latter configuration consists of two sets of live electrons each having parallel spin that will contribute a total 20K towards the exchange energy. Therefore, there is a total gain of 4K in exchange energy if copper consists of the configuration 3d104sl. Though, in accomplishing this configuration, there will again be a loss in the energy in promoting an electron from 4s orbital to 3d orbital, which happens to be less than 4K, the gain in exchange energy. Therefore, the 3d104s1 configuration becomes more stable than 3d94s2.
It is as well worth mentioning here that although the 4s orbitals are occupied before 3d orbitals, we can't state that they are always more stable. However, the ionization of the transition elements occurs by the loss of ns electrons first. What happens in reality is that if the electron is ionized from any transition element, state the one from 3d series, the efficient nuclear charge experienced via the 3d electrons is greatly-improved over that of any 4s electron as a direct effect of the greater stability acquired by the 3d orbitals in the due course of filling. As a result, the 3d orbitals are expected to drop significantly in energy beneath the 4s orbital. Therefore, ionization of two or more electrons from an atom of a transition element will occur with the elimination of 's' electrons in preference to the 'd' electrons.
Therefore, we notice that it is the net effect of all the forces, comprising nuclear-electronic attraction, shielding of one electron via others from the nuclear charge, inter-electronic repulsion and exchange forces that finds out the stability of electronic configuration.
The transition elements have some common properties, which are represented below:
1) All the elements are metals and form alloys by one another and with other metallic elements.
2) They are strong, hard, ductile, malleable, high melting and high boiling. They are very good conductors of heat and electricity.
3) Most of them are sufficiently electropositive to dissolve in mineral acids however a few are noble - that is, they encompass such low electrode potentials that they are not affected by simple acids.
4) They generally show multiple oxidation states.
5) They form coordination compounds or ions. However, the chemistry of the transition elements is mostly related by the use of d and also 5 and p orbitals in making coordination compounds.
6) The transition metal complexes are generally colored.
7) Most of the compounds of transition elements are paramagnetic.
8) Most of these elements and their compounds act as the catalysts for chemical reactions.
Periodic trends in properties:
We have studied the significant properties of transition metals in general. As we are familiar that the transition metals are an integral part of the periodic table, similar to the main group elements, the transition metals are as well expected to show periodicity in their properties.
Some of the significant properties of the elements of 3d - series are listed in the table shown below. If we observe the data in the table carefully, we will notice that all along a period, these properties differ much less from one element to the other as compared to the main group elements. However, the horizontal similarity amongst the d-block elements is well marked, yet the chemistry of the elements of first transition series varies considerably from that of the elements of the second and third transition series, which are incidentally more identical to each other. This difference in trends in the properties of d-block elements from those of 's' and 'p' block elements occurs from a fundamental difference in their electronic configuration.
As in building up of elements from lithium to fluorine, the electrons are added to the outermost shell, in the case of transition metals, the electrons are added to inner (n-1)d sub-shell.
Table: Properties of 3d elements
Atomic Radii, Atomic Volume and Density:
From the table above, we can notice that there is a gradual decrease in the atomic radius across a row of transition elements. On passing from left to right, extra positive charges are placed on the nucleus and respectively electrons are added to the (n-1)d orbitals. As the electrons in the d orbitals shield the ns electrons and as well themselves from the nuclear charge incompletely, effective nuclear charge felt by them increases and therefore a contraction in size takes place.
Though, it is significant to emphasize here that shielding of the outer ns electron(s) by (n-l)d electron(s) is more efficient than the shielding of an ns electron by the other ns electron (or that of an np electron via another np electron). This is why the decrease in atomic radius from sodium to chlorine is more than that from scandium to copper. The elements that take place instantly after the transition elements are smaller than expected from the simple extrapolation from the group elements. This is because the cumulative effect of incomplete shielding given by (n-1)d10 electrons and thus, the effective nuclear charge fell by the outer electrons of the elements from gallium to krypton is more than that if the d-orbitals had not been slowly filled in the transition elements.
Though, it is significant to emphasize here that shielding of the outer ns electron(s) by (n-l)d electron(s) is more efficient than the shielding of an ns electron by the other ns electron (or that of an np electron via another np electron). This is why the decrease in atomic radius from sodium to chlorine is more than that from scandium to copper. The elements that take place instantly after the transition elements are smaller than expected from the simple extrapolation from group elements. This is because the cumulative effect of incomplete shielding given by (n-1)d10 electrons and thus, the effective nuclear charge fell by the outer electrons of the elements from gallium to krypton is more than that if the d-orbitals had not been slowly filled in the transition elements.
The group trends in atomic radii of the transition elements are parallel to those noticed in the 's' and p-block elements. As we go down the group, there is an increase in the atomic size up to second transition series. This is not surprising in view of the fact that electrons enter the 4d orbital in the second transition series. Though, the size of elements of third transition series is nearly identical to that of the elements of second transition series as the filling of 4 orbitals in the lanthanides.
The atomic volume of an element is directly associated to its size and, thus, atomic volumes follow the similar trend as the atomic size. Likewise density is as well associated to the size of the element. The smaller the size, the higher is the density of the element. Therefore there is a general trend of increasing density across the elements of a transition series. For 4d and 5d elements, this increase is not that regular as the increase in densities for the 3d elements. All along the group as well, the density increases. The increase in density in the d block groups is more than that in the 5 and p block groups.
Melting and Boiling Points:
The boiling and melting points of the transition elements are generally high. The melting points of the elements based on the strength of the metallic bond. As we are familiar, that the transition metals crystallize in the metallic lattices. The strength of metallic bond rises with the availability of the electrons to participate in the bonding via delocalization. It will observe that between calcium and scandium (where d electron first appears), there is a jump of almost 700 K in the melting point. The presence of one or more unpaired d electrons therefore leads to higher interatomic forces and thus, high melting and boiling temperatures. Therefore, we can think that by the increasing availability of the unpaired d electrons, the strength of metallic bond rises, resulting in the higher melting points. However we can't generalize the argument because whenever we move across any period in the periodic table, the melting point raises upto the middle of each transition series and then it reduces with the starting of electron pairing for the elements of first transition series.
Fig: Melting points of alkali, alkaline earth and transition metals
There is a sharp decrease of melting point at manganese, which consists of five unpaired d electrons. Though, the softness and low melting point of Zn, Cd and Hg (note that, Hg is a liquid) in which all the electrons are paired up can tentatively be described on the above basis. The melting points of elements of the first transition series are comparatively lower than those of the elements of the second and third transition series. This trend is very well described in the figure above.
The periodic trends in the boiling points are alike to those in the melting points. As the method of boiling needs almost complete breaking of bonds and such metallic bonding exists in the liquid state to certain extent, high temperatures are essential. Thus, the boiling points of the metals are much higher than their melting points.
In case of transition metals as well, the variation of ionization energy across the periods and down the groups parallels quite closely to the trend in atomic size. This is illustrated nicely in the figure shown below.
As we move across a period, the effective nuclear charge experienced by ns2 electrons goes on rising causing the shells to shrink in size and therefore making it hard to remove the electrons. Therefore all along a period, the ionization energy increases. The second and third ionization energies follow the similar pattern, apart from for the second ionization energies of Cr and Cu that are comparatively higher due to the additional stability of 3d5 and 3d10 configurations. The ionization energies of the elements of second and third transition series as well follow the similar trend all along the period. As the decrease in size of the transition metals is less than that of the main group elements all along a period, the ionization energies tend to increase all along the series just slightly as compared to the main group elements. As s and d electrons don't differ much in energy, the difference in the successive ionization energies is comparatively small.
Fig: Ionization energies of second and third transition series
As we move down the group from elements of first transition series to those of the second, there is a decrease in the ionization energy. However it again rises when we move further down the group from second to the third transition series. This trend is consistent with relatively small size of the atoms of elements of the third transition series. This is because the insertion of lanthanides that causes the third row transition elements to encompass greater than expected effective nuclear charge.
Transition elements contain fairly low values of electronegativity. It rises from Sc to Cu by a fall at Mn and Zn. Though, this increase in electronegativity is much slower as the additional electron is being added to the inner shell that gives relatively good shielding to the outer electrons from the nucleus. The increasing electronegativity from Sc to Cu signifies that the elements become slightly less metallic and this is reflected in the increasing positive electrode potentials of their ions M2+ and M3+
Prior to going into the details of the variation in the electrode potential of the transition elements, let us first illustrate the concept of electrode potential. Whenever a metal is put in a solution of its ions a potential difference is set up between the metal and the solution. There is a tendency of the metal ions to leave the metal lattice and go to the solution therefore leaving a surplus of electrons and therefore a negative charge on the metal; there is as well a reverse tendency for the metal ions from the solution to deposit on the metal leading to the positive charge on the metal. In practice one of these effects is more than the other, bringing about a potential difference between the metal and solution. The value of this potential difference for a specific metal based on the nature of metal, the concentration of metal ions in solution and the temperature. Via convention, the potential difference set up in a 1M solution of metal ions at 298K is known as the standard electrode potential. This is not possible to measure the standard electrode potentials absolutely. Standard electrode potentials, thus, have to be measured against certain reference standard; the one adopted is the hydrogen electrode. This comprises of hydrogen gas at one atmosphere pressure in contact by a 1M solution of its ions at 298 K.
In common, we can state that more negative the value of the electrode potential for the couple Mn+/M, more is the reducing power of the element. Likewise, more positive is the value of electrode potential for the couple Mn+/M, more is the oxidising power of the element. Electrode potential is the measurement of electropositive character and the reactivity of the metals. In common all along a period, there is a decrease in electropositive character. The reactivity of metals also decreases along a period and down a group.
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