Spectrochemical Series of some d-block Transition Metal Complexes

Spectrochemical Series of some d-block Transition Metal Complexes (Adapted from: Inorganic Chemistry: Discovery Laboratory Experiments for Part 1 by Gary Wulfsberg) Introduction: Description of some of the early experiments and theories that developed toward the end of 1800 s had led scientists in particular, the chemists in the study of coordination compounds. In modern terms, there are two useful concepts that led to described them; these are the central metal ion (atom) and its ligands. The transition elements, particularly form ions which in certain medium can exist independently and which in other environments can act as central ions that can combine with ions or molecules to form complex species or coordinated compounds. A good example is Fe(III): It occurs as Fe 3+, but also as FeCl 2 +, FeCl 4 -, Fe(CN) 6 3- and Fe(H 2 O) 6 3+, to list only a few of its many complex ions. In each of these complex ions, the iron is called the central metal ion, and the groups that combine the central metal ion are called to be the ligands. A ligand may be an ion or a polar molecule. It is a Lewis base that can donate a pair of electrons that functions as a covalent bond between the central ion and the ligand. When both electrons are supplied by the bound species, the bond is called a coordinate covalent bond. Typical ligands are ions of non-metals or ions or molecules that contain a non-metal atom, which in turn has an electron pair to donate to the central transition metal ion. In two of the complex ions mentioned, the Cl - is the Lewis base, and in the water complex, water s oxygen acts as the electron donor, the cyanide ligand, is probably bound through the electron pair of carbon, but in some cases, the pair of electron on nitrogen will serve as the coordinate bond for Fe(lll). The number of ligands attached to the central ion is usually termed as coordinate number. In principle, positive ions of every metal in the periodic table accept electron density to some degree and can therefore coordinate surrounding electron donors, even only if weakly. As an example, the solvation of the surrounding H 2 O molecules with Na + is a weak coordination. The ability to make fairly strong, directional bonds by accepting electron pair from neighbouring molecules or ions is a characteristic of a transition metal element. Thermodynamically speaking, coordination occupies the boundary between weak intermolecular interactions and stronger covalent and ionic bonding. Thus, the heating of an aqueous green Nickel(ll) nitrate solution disrupts the Ni-H 2 O bonds at the temperatures well below those required to break the covalent bond in the NO 3 -. Based from statistical mechanical considerations, the enthalpy change in breaking a +2 transition metal ion away from the coordinated water molecule falls in the range of 170-210 kj/mol. This is far less than the strongest chemical bonds, but it is by no means small. Metal ions with +3 charge have still stronger coordinate covalent bonds with water. Of particular interest in the field of coordination chemistry are the magnetic properties, spectral properties, energetics and colours of the coordination compounds. These properties are variable and 1

striking, and the rich variety of geometrical arrangements in which ligands position themselves around metal atoms. From the results of these studies, it is possible to outline the some of the properties, and then apply the present modern bonding theories to account and explain the observed physico-chemical properties of the complex being studied. One of the most aesthetically valued properties of the coordinated complexes is their hue or colour. Colours arises because complexes often absorbs light fairly effectively in some portion of the visible spectrum. The colour perceived in the sample is the colour complementary to that which most strongly absorbed. To illustrate, the famous deep blue complex [Cu(NH 3 ) 4 ] 2+ absorbs strongly roughly in the 620nm region of the visible spectrum. This portion corresponds to the orange region, which in turn turns out to be the complementary colour of blue. The colour wheel below shows that the colour opposite to the wheel is the complement of the colour being studied. Figure 1: A typical colour wheel Absorption in the visible part of the spectrum is due to the transitions from lower to higher energy electronic states. The colour of the coordination complex is determined by the transition of the electrons in the neighbourhood of the central metal atom, because the presence of such atoms is required for the compound to be coloured. As a student of science, it is important to understand the origins of the absorptions that give rise to colours of the coordination compounds, and why the colour changes as they do when the ligands are changed. Another interesting property of a substance is its behaviour when subjected to an external magnetic field. Substances can be classified as paramagnetic or diamagnetic, according to whether or not they are attracted into a magnetic field. An experiment to demonstrate the universal susceptibility of substances to the influence of the magnetic field is shown below. 2

Figure 2: Schematic of a typical magnetic susceptibility apparatus. A sample in the form of the cylinder is suspended so that its bottom is between the poles of a powerful magnet but its top extends out of the field. It is weighted very accurately and then reweighted when the magnet is moved away. When the prototype study is carried out, it was found out that if the net force in the sample changes in the presence of the magnetic field. Substances that are repelled by a nonuniform magnetic field weight less when dipped into one and are diamagnetic; substances that are attracted by a magnetic field weighted more and are paramagnetic in character. The weighings just described give numerical values for the magnetic susceptibility of a diamagnet and is negative and small, while the susceptibility of a paramagnet is positive and can be quite large. Recall from molecular orbitals and spectroscopic results that paramagnetism is associated with atoms, ions or molecules that contain one or more electrons with unpaired spins. Diamagnetic substances have the spins of all electrons paired up. In many cases, the number of unpaired electrons per molecule can be counted, based on the magnitude of its magnetic susceptibility. On a molar basis, a substance with two unpaired electrons per molecule is pulled into a magnetic field more strongly than a substance with only one unpaired electron per molecule, but is less strongly with three unpaired electrons per molecule. These facts come up in connection with coordination complexes, whereas the great majority of other substances are diamagnetic. Among complexes of a given metal ion, the number of unpaired electrons, as observed by the magnetic susceptibility, varies with the identity of the ligand. As an example, both [Co(NH 3 ) 6 ] 3+ and [Co(ox) 3 ] 3- have the similar molecular geometry, that is an octahedral like structure. However, the former is diamagnetic and the latter is paramagnetic. On the molecular level, the magnetic properties and colour of the transition metal complexes have a common source, one that involves the d electrons on the central transition metal and the way they are perturbed by the approach of the ligands. To further develop the idea mentioned above, it is imperative that a discussion on the bonding theory of coordination compounds being mentioned. The simplest model of a chemical bond that can be applied to coordinated compounds is the Valence Bond theory. One of its highlights is the orbital occupancy of the central metal ion in the complex. Also, it is able to predict the magnetic properties of the complex by forming either an inner orbital complex or an outer orbital complex. By just simply knowing the magnetic susceptibility of the 3

complex being studied, Valence Bond can be used to elucidate its geometric structure. A square planar complex may be distinguished from the tetrahedral complex that have that same colour by just looking at the free ion configuration of the central ion and its orbital occupation that includes the hybridization of the orbitals in the outer shell. Unfortunately, one of the limitations of the Valence Bond model is that the theory could not account for the colour of the complex. A second theory of bonding in transition metal complexes, that has been extensively applied over the past 50 years, is the Crystal Field Theory (CFT). It differs from Valence Bond theory in that it views the complex as held by purely electrostatic interactions, that is in its simple form, CFT ignores covalent bonding. The most significant aspect of the theory is that, it is concern with the effect of the ligands have on the energies of the d-orbitals. To illustrate CFT, consider the five d-orbitals of the central transition metal in a hypothetical complex. Suppose that the complex would be formed is an octahedral one. From quantum mechanics, one could see that if the ligands are very far from the central metal, the d-orbitals would be quintuply degenerate. The figure below shows the structure of the five d-orbitals. Figure 3: The quintuply degenerate d-orbitals. Since the d-orbitals have a sense of directionality, then the approach of the ligand along the Cartesian axes would cause different degrees of electrostatic interaction with the orbitals. By examining the figure above, the d-orbitals that lie along the coordinate axes will be highly destabilized by the approach of the ligand in comparison to those that do lie on between the axes. Thus, the two d-orbitals would be on a higher energy in comparison with the remaining three d-orbitals. 4

Figure 4: The splitting of the d-orbitals in the approach of the ligands in the quintuply degenerate d-orbitals. For an octahedral system, the splitting of the five d-orbitals results in the formation of a triply degenerate and doubly degenerate quantum state. The energy difference between the splitting of the d- orbitals is termed as the Crystal Field Splitting Energy (CFSE). Based from the study conducted by R. Tsuchida, the magnitude of the CFSE varies with the intensity of the ligand. From such observations, he had proposed that the ligands can be arranged in a spectrochemical series. The ligands are arranged in such a way that the CFSE increases. A typical spectrochemical series is shown below: I - < Br - < S 2- < SCN - < Cl - < NO 3 - < F - < OH - < C 2 O 4 2- < H 2 O < NCS - < CH 3 CN < NH 3 < en < bipy < phen < NO 2 - < PPh 3 < CN - < CO Moreover, the CFSE is a conjugal property; it also depends on the characteristics of the central metal. From empirical observations, the way the central metal affects the CFSE are as follows: 1. As the Oxidation state of the central metal increases, CFSE increases. 2. For a particular group of metals, CFSE increases as the metal goes down the group. Shriver and Atkins had explained the trends as follows: The variation with oxidation number reflects the smaller size of more highly charged ions and the consequent smaller metal-ligand distances. The second factor reflects the improved metal; ligand bonding of the more expanded 4d and 5d orbitals compared with the compact 3d orbitals. Calculation of CFSE is overall difficult except in simple cases where there is just a single electron. However, in principle, the CFSE can be calculated by knowing the wavelengths at which the electronic 5

transitions occur. This can be done by using the appropriate Tanabe-Sugano diagram. In a practical sense, based from Atkins, Tanabe-Sugano diagrams are correlation diagrams that depict the energies of the electronic states of complexes as a function of the strength of the ligand field. A typical Tanabe- Sugano Diagram is shown below. Figure 5: A typical Tanabe-Sugano diagram for a d 2 complex. One of the major success of CFT is that it is able to account the colour that the complex exhibits. According to CFT, if the energy of the light that strikes the complex is equal to the CFSE, the light can be absorbed and the electron can be raised from the lower set of d-orbitals to the higher set. The energy of this absorbed light of course depends on the magnitude of CFSE. For most complexes, the magnitude of CFSE is such that the frequencies of light that are absorbed reside in the visible spectrum. Since the colour of the light is related to the frequency, the colour of the complex depends on the frequencies that are absorbed when white light is reflected from it or passes through it. In other words, the observer sees the complement of the colour of the light that is absorbed. On the other hand, not all d-block elements in the periodic table do exhibit the properties of the transition metal elements. One example is Zinc, whose electronic configuration is [Ar] 3d 10 4s 2. Since the d-orbitals are fully filled then from theoretical considerations, it can be conjectured that it is highly stable since the exchange energy would be at its maximum. Also, it could be hypothesized that the energy required for the promotion of the electrons in the d-orbitals of zinc is very large in such a way that it is outside the visible region of the electromagnetic spectrum that is why; its promotion is not seen when exposed in the visible light. Also, based from the canonical ensemble and the generation of the Boltzman distribution, which states that the probability of finding a particle in the high energy 6

quantum mechanical state is very small. That is why, even though the transition is possible, the probability of such event is very minimal. Thus, Zn(ll) appears to be colourless in aqueous environment. In this experiment, a spectrochemical series of the ligands will be constructed. We will begin with the visual eyeball approximations followed by quantitative measurements by means of a UV-Vis spectrophotometer. The success of the experiment depends on how will you analyse your results. Procedure: Part A: Preparation of Complexes and Their ordering by Visual Inspection 1. Prepare seven clean and dry 4-inches test tubes. To each test tube, place 2 ml of the transition metal ion assigned to your group. Label your test tubes A-G. 2. Add the following reagents to their corresponding test tubes stopper them and mix. Record your observations. A. 2 ml of distilled H 2 O B. 2 ml of 1.0 M Sodium Glycinate, Na + NH 2 CH 2 CO 2 - C. 2 ml of saturated K 2 C 2 O 4 D. 2 ml of 1.0 M NH 3, buffered E. 2 ml of 1.0 M pyridine F. 2 ml of 1.0 M Sodium ethylenediaminetetraacetate (EDTA) G. 2 ml of 1.0 M ethylenediamine, buffered 3. Does the ligand affect the colour of the transition metal complex? Explain. 4. Arrange the colour of the test tubes in rainbow order. List the ligands in order, start your list with the end that is closest to the colour of the transition metal ion in the pure water test tube. 5. Is the ligand ordering for Ni(II) complexes are reasonably consistent with the ordering of the ligands for the complexes of Co(II) and Cu(II)? As a group, check the possibility that making slight rearrangements would make a more nearly metal-independent listing of the ligands in order. Your final group list is a qualitative spectrochemical series of ligands. Part B: Spectral Analysis of the Transition Metal Complexes 1. Obtain the spectra of all the complexes. Measure the spectra in the visible and near-ir region using a spectrophotometer. 2. For each complex, tabulate the corresponding wavelengths of the peaks. Convert the resulting wavelengths to wavenumbers (in cm -1 ). 3. In the Cu(II) complexes you will observe only one d-to-d transition. The energy of this transition is the octahedral ligand field splitting, Δ o. In other octahedral complexes, you will observe more than one d-to-d transition; the peak that has the longest wavelength in these complexes is taken as Δ o. Record the Δ o for each of your complexes. 4. The Δ o can be approximately factored out into two components, an f factor due to the ligand and a g factor due to the metal 7

(Eq. 1) By convention, we can assign an f factor of 1.000 to the water molecule as the ligand. From the Δ o value of your hydrated metal ion, compute the g value for your metal ions. Tabulate them. 5. Your experimental g factor can then be used to calculate the f factors for the other ligands in your complexes. To obtain the f factor, we note that (Eq. 2) Dividing Eq. 2 with Eq.1 will cancel the g factor and yields (Eq. 3) From this equation, calculate the f factors of your ligands. 6. List your ligands in order of increasing order of their f values. Compare the resulting list with the one that you obtained by eyeball inspection in Part A. 7. Analyze the periodicity of your spectrochemical series: can you note any grouping of nitrogenand oxygen-donor ligands in the spectrochemical series? REFERENCES [1] Atkins, P.W. Inorganic Chemistry. Third ed.oxford University Press. Oxford 1999. [2] Barrow, G.M. Physical Chemistry. 6 th ed. Mc Graw Hill. USA 1996. [3] Cotton, F.A., Geoffrey Wilkinson, and Paul L. Gaus. Basic Inorganic Chemsitry. Third ed. John and Wiley. New York 1995. [4] Housecroft, C.E. and Alan G. Sharpe. Inorganic chemistry. 2 nd ed. Pearson Prentice Hall. New Jersey. USA 2005. [5] Huheey, James E. Inorganic chemistry: principles structure and reactivity. 3rd ed. Cambridge : Harper and Row, Publishers, 1983 8