Initials and surname: M Letlhabi
Student number: 33515859
Module code: NCHE321
Title:
Electronic absorption of some Cu(II) complexes.
Abstract:
When reacting with water and ammonia solution, copper (II) ion creates an octahedral complex,
which is the most frequent kind of complex. Cu(NH3)n(H2O)6-n2+, which is the formula for this
complex. The molecules of water and ammonia are the in this instance, ligands. In this experiment,
we deterimined the maximum wavelength absorbance using the visible spectra between the range
900nm to 500nm for each complex with a scale of 0–1. The 1M NH3 recorded relatively low values
for Δ° as there were no other ligands added to the complex, on the contrary the crystal splitting
energy values of glycine and ethylenediamine were relatively small. This experiment aims to record
electronic absorption spectra for some Cu(II ) complexes and to determine crystal field splitting
energies thereof.
Objectives:
•
•
•
To record electronic absorption spectra for complexes of Cu(II)
To study the effect of ligands on the Crystal field splitting energy (DE)
To determine crystal field splitting energies
Introduction:
Transitional metals have partially filled 3d orbitals. Because of the incompleteness of these orbitals,
they have numerous traits that are unique to them. One of these abilities is the ability to generate
coloured complexes with ligands. Electrons of transition metals tend to become tremendously
excited and unstable during their reaction with the ligands. The subsequent absorption and release
of energy in the form of light is synonymous with electron excitation and de-excitation (Griffith,
1964). This absorption occurs in the visible part of the electromagnetic spectrum, hence the
compounds have a colourful trait. Light in this part of the spectrum yields the promotion of delectrons from a lower to a higher energy level. The resulting spectra are commonly referred to as
Electronic Spectra (Shimada and colleagues, 2011). As a result of the magnitude of the quanta at
stake, electronic molecule transitions are always accompanied by vibrational and rotational energy.
Resulting in the detection of a band spectrum. There are several assumptions used to relate electrical
absorption spectra to the structure of an electronic device. Crystal Field Theory (CFT) is one of those
renowned assumptions. In this method, it is believed that the hypothetical interaction between a
metal particle and the ligands is electrostatic (NCHE321 Practical Guide, 2023, p.51). The
deteriorated metal d orbitals are divided into sub-groups in the presence of the electrostatic field
formed by the ligands. For instance, in an octahedral field the d orbitals split into two sets, the inter-
axial set (dyz, dxy, dxz) symbolized by t2g and the orthogonal set and the hub set of higher energy
dz2 and dx2-y2.
When copper (II) ion reacts with water and ammonia solution, it produces an octahedral complex
which is the most frequent type of complex. The formula of the complex formed is Cu(NH3)n(H2O)62+
n . In this instance, water and ammonia molecules are the ligands. Six ligands (H2O and NH3) are
present in this case of the octahedral complex. In this case of the octahedral complex, six ligands of
water and ammonia form an octahedron around the metal ion (Figgs and Hitchman, 2000). The
crystal splitting energy is therefore the difference in energy between the t2g and eg orbitals in an
octahedral system. There’s a variance between the magnitude of the splitting of the t2g and eg
orbitals of one octahedral complex to another. This is dependent on the identity of the metal ion, the
charge on this ion and its oxidation state, the nature of the ligands coordinated to the meal ion,
whether the metal is a 3d, 4d or 5d element.
Experimental procedure:
For n=0, 1 mol dm-3 of Cu(NO3)2 stock solution was diluted using a pipette 50.0 cm3 in a volumetric
flask. 5.0 cm3 of the 1 mol dm-3 Cu(NO3)2 was pipetted into a small beaker, thereafter 14g of solid
ammonium nitrate was added. The solution was stirred and warmed to about 300C using a magnetic
stirrer. 5 cm3 of 1 mol dm-3 NH3 was added and the stirring continued until the temperature reached
about 300°C again. For n=1, 1.0 cm3 of the solution was diluted to 25 cm3 with 2 mol dm-3 NH4NO3
solution. For n=2 and n=3 the same method was used as outlined for n=1 but replaced 1 mol.dm-3
NH3 with 2 and 3 mol dm-3 NH3 respectively. For n=4, 1 cm3 of concentrated NH3 was added to 1 cm3
of the 1 mol.dm-3 Cu(NO3)2 and the solution was diluted with water to 50 cm3. 1.0 cm3 of the
mol.dm-3 Cu(NO3)2 solution and 5.0 cm3 of the 0.4 mol dm-3 ethylenediamine solution was pipetted
into a 50 cm3 volumetric flask and made to the mark with distilled water. 1.0 cm3 of the 1 mol.dm-3
Cu(NO3)2 solution was pipetted into a 50 cm3 volumetric flask and solid glycine was added to obtain a
concentration of 0.40 mol.dm-3. 20cm3 of 0.1 mol.dm-3 NaOH was added and made to the mark with
distilled water. The wavelength of maximum absorbance for each complex was determined using the
visible spectra over the ranges of 900 to 500 nm with absorbance scale of 0 to 1.
Results And Discussion:
Mass calculation of Glycine:
As observed from the graph , the highest energies were recorded at values of n=0 for each of the
procedures involving 2M and 3M NH3. The 1M NH3 was the only one to record low values for Δ° since
there were no other ligands added to the complex, The values of the crystal splitting energy of
glycine and ethylenediamine were relatively small and this can be attributed to their nature. As
organic compounds they are hardly capable of forming co-ordinate bonds with other negatively
charged ligands. In this case their crystal splitting energy depends mostly on their nature which is an
indication of their stability. The energy associated with the ethylenediamine was less than that of
glycine so it can be said that ethylenediamine is less stable than glycine. A comparative analysis of
the complexes revealed that ligand substitution influenced the positions and intensities of the
absorption bands. The presence of different ligands led to variations in the ligand field strength,
resulting in spectral shifts (Figgs and Hitchman, 2000). The λmax values suggest that the electronic
transitions in these complexes involve different energy levels. The variation in λmax can be attributed
to the nature of ligand-field splitting, which is influenced by the ligand's electronic properties and
coordination geometry. Ligands with different electron donating abilities and sizes lead to variations
in the energy levels of d-orbitals in the central copper atom (Ayscough, 1967). The observed trend in
the graph aligns with the spectrochemical series, which predicts ligand-field strengths based on
ligand properties. In this experiment, ammonia (a stronger ligand) caused a larger shift in λmax
compared to ethylenediamine which isna weaker ligand, resulted in the smallest shift. This trend is
consistent with the ligands' ability to interact with the copper atom and induce ligand-field splitting.
Conclusion:
In this study, the electronic absorption spectra of Cu(II) complexes were investigated, highlighting the
effects of ligand substitution on the observed absorption bands. It can be concluded that the crystal
splitting energy of some copper complexes does increase due to increasing concentration and
number of NH3 ligands present in the complex. Different ligands generate crystal fields of different
strengths, hence different colors were observed. The spectral shifts were attributed to ligand field
transitions and electronic excitations, which are influenced by ligand properties and coordination
geometry. This experiment provided valuable insights into the electronic structure of transition metal
complexes and their sensitivity to ligand variations.
References:
[1] ”Ligand Field Theory and Its Applications”,Wiley-VCH, New York,(2000). B.N.Figgs,M.A.Hitchman
[2] A.Lund, M.Shiotani, S.Shimada,“Principles and Applications of ESR Spectroscopy”,Springer New
York (2011).
[3] C.J.Ballahausen, “Introduction to Ligand Field Theory”,McGraw Hill Book Co.,NewYork (1962).
[4] P.B. Ayscough,”Electron Spin Resonance in Chemistry”, Mathuen Co.,Ltd.,London (1967).
[5] R.L.Carlin, “Transition Metal Chemistry”, Marcel Dekker,New York (1969).
[6] Journal of “Coordination Chemistry Reviews”.
[7] Journal of Spectrochimica Acta A Elsvier.
[8] J.S.Griffith, “Theory of Transition Metal Ions”, Cambridge University Press,Oxford (1964). [9] NCHE
321 Practical Guide 2023, NW