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2. Submitted to:Dr. Ali MohsinSubmitted by:Adina Tatheer10060607-050BS Chemistry (V)Section A 3. Q no. 1: Describe the magnetic properties of complex.Magnetic properties can be determined by looking at a compounds electron configurationand the size of its atoms. Since Magnetism is created by the spin of electrons, we canlook at how many unpaired electrons are present in a specific compound and determinehow magnetic the compound is. For this purpose we will be evaluating the d-blockelements or Transition Metals* (TMs) because they tend to have a large number ofunpaired electrons.Introduction:The magnetism discussed in this article is paramagnetism. Paramagnetism occurs whenthere are one or more unpaired electrons in a compound. (The opposite, when allelectrons are paired, is called diamagnetism). Di- and para-magnetism are often affectedby the presence of coordination complexes, which the transition metals (d-block) readilyform.Singular electrons have a spin, denoted by the quantum number ms as +(1/2) or (1/2).This spin is negated when the electron is paired with another, but creates a slightmagnetic field when the electron is unpaired. The more unpaired electrons, the morelikely paramagnetic a material is. The electron configuration of the transition metals (d-block) changes when in a compound. This is due to the repulsive forces betweenelectrons in the ligands and electrons in the compound. Depending on the strength of theligand, the compound may become paramagnetic or diamagnetic.Ferromagnetism:Some paramagnetic compounds are capable of becoming ferromagnetic. This means thatthe compound shows permanent magnetic properties rather than exhibiting them only inthe presence of a magnetic field. In a ferromagnetic element, electrons of atoms aregrouped into domains, where each domain has the same charge. In the presence of amagnetic field, these domains line up so that charges are parallel throughout the entirecompound. Whether a compound can be ferromagnetic or not depends on how manyunpaired electrons it has and on its atomic size. 4. Small atoms pair up too easily and their charges cancel. Large atoms are difficult to keep together, their charge interaction is too weak.Therefore, only the right sized atoms will work together to group themselves intodomains. Elements with the right size include: Fe, Co, Ni. That means that Fe, Co and Niare paramagnetic with the capability of permanent magnetism; they are alsoferromagnetic.Ligand Field Theory BackgroundAn element can have up to 10 d electrons in 5 d-orbitals, dxy, dxz, dyz, dz2, and dx2-y2. During the formation of a complex, the degeneracy (equal energy) of these orbitals isbroken and the orbitals are at different energy levels.(Assuming a 6-ligand compound)In an octahedral complex, the ligands approach along the x, y, and z axes, so therepulsion is strongest in the orbitals along these axes (dz2 and dx2-y2). As a result, thedz2 and dx2-y2 orbitals are higher in energy than the dxy, dxz, and dyz orbitals. In a tetrahedral 5. complex, the splitting is opposite, with the dxy, dxz, and dyz orbitals higher in energy toavoid the ligands approaching between the axes. The splitting in a square planar complexhas four levels (lowest to highest): dyz and dxz, dxy, dz2, dx2-y2.Depending on the strength of the ligand, the splitting energy between the different d-orbitals may be large or small. Ligands producing a smaller splitting energy are calledweak field ligands, and those with a larger splitting energy are called strong fieldligands.Filling of d-orbitals in a complex:Hunds Rule states that electrons will fill all available orbitals with single electrons beforepairing up, while maintaining parallel spins (paired electrons have opposing spins). For aset of degenerated d-orbitals (not in a complex), electrons fill all orbitals before pairing toconserve the pairing energy, otherwise needed. With the addition of ligands, the situationbecomes more complicated. The splitting energy between the d-orbitals means thatadditional energy is required to place single electrons into the higher-energy orbitals.Once the lower-energy orbitals have been half-filled (one electron per orbital), an 6. electron can either be placed in a higher-energy orbital (preserving Hunds rule) or pairup with an electron in a lower-energy orbital (when the splitting energy is greater than thepairing energy). The strength of the ligands determine which option is chosen.With a strong-field ligand, the splitting energy is very large and low-spin complexes areusually formed. With a weak-field ligand, the electrons can easily enter the higher-energyorbitals before pairing (high-spin).How does this relate to magnetism?Low-spin complexes contain more paired electrons since the splitting energy is largerthan the pairing energy. These complexes, such as [Fe(CN)6]3-, are more oftendiamagnetic or weakly paramagnetic. High-spin complexes usually contain moreunpaired electrons since the pairing energy is larger than the splitting energy. With moreunpaired electrons, high-spin complexes are often paramagnetic.The unpaired electrons in paramagnetic compounds create tiny magnetic fields, similar tothe domains in ferromagnetic materials. The higher the number of unpaired electrons(often the higher-spin the complex), the stronger the paramagnetism of a coordinationcomplex. We can predict paramagnetiism and its relative strength by determiningwhether a compound is a weak field ligand or a strong field ligand. Once we havedetermined whether a compound has a weak or a strong ligand, we can predict itsmagnetic properties: 7. Q no. 2: Explain the CF in the octahedral and tetrahedral symmetry alsoexplain the factors affecting the magnitude of .CF theory:CF theory tried to describe the effect of the electrical fieldof neighboring ions on the energies of the valence orbitals of an ion in a crystal. Crystalfield theory was developed by considering two compounds: manganese(II) oxide, MnO,and copper(I) chloride, CuCl.Octahedral Crystal Fields:Each Mn2+ ion in manganese(II) oxide is surrounded by six O2- ions arranged toward thecorners of an octahedron, as shown in the figure below. MnO is therefore a model foran octahedral complex in which a transition-metal ion is coordinated to six ligands.What happens to the energies of the 4s and 4p orbitals on an Mn2+ ion when this ion isburied in an MnO crystal? Repulsion between electrons that might be added to theseorbitals and the electrons on the six O2- ions that surround the metal ion in MnO increasethe energies of these orbitals. The three 4p orbitals are still degenerate, however. Theseorbitals still have the same energy because each 4p orbital points toward two O2- ions atthe corners of the octahedron.Repulsion between electrons on the O2- ions and electrons in the 3d orbitals on the metalion in MnO also increases the energy of these orbitals. But the five 3d orbitals on theMn2+ ion are no longer degenerate. Lets assume that the six O2-ions that surround eachMn2+ ion define an XYZ coordinate system. Two of the 3d orbitals (3dx2-y2 and 3dz2) onthe Mn2+ion point directly toward the six O2- ions, as shown in the figure below. Theother three orbitals (3dxy, 3dxz, and 3dyz) lie between the O2- ions. 8. The energy of the five 3d orbitals increases when the six O2- ions are brought close to theMn2+ ion. However, the energy of two of these orbitals (3dx2-y2 and 3dz2) increases muchmore than the energy of the other three (3dxy, 3dxz, and 3dyz), as shown in the figurebelow. The crystal field of the six O2- ions in MnO therefore splits the degeneracy of thefive 3d orbitals. Three of these orbitals are now lower in energy than the other two.By convention, the dxy, dxz, and dyz orbitals in an octahedral complex are calledthe t2g orbitals. The dx2-y2 and dz2 orbitals, on the other hand, are called the eg orbitals. 9. The easiest way to remember this convention is to note that there are three orbitals inthe t2g set.t2g: dxy, dxz, and dyz eg: dx2-y2 and dz2The difference between the energies of the t2g and eg orbitals in an octahedral complex isrepresented by the symbol o. This splitting of the energy of the d orbitals is nottrivial; o for the Ti(H2O)63+ ion, for example, is 242 kJ/mol.The magnitude of the splitting of the t2g and eg orbitals changes from one octahedralcomplex to another. It depends on the identity of the metal ion, the charge on this ion, andthe nature of the ligands coordinated to the metal ion.Tetrahedral Crystal Fields:Each Cu+ ion in copper(I) chloride is surrounded by four Cl- ions arranged toward thecorners of a tetrahedron, as shown in the figure below. CuCl is therefore a model fora tetrahedral complex in which a transition-metal ion is coordinated to four ligands.Once again, the negative ions in the crystal split the energy of the d atomic orbitals on thetransition-metal ion. The tetrahedral crystal field splits these orbitals into thesame t2g and eg sets of orbitals as does the octahedral crystal field.t2g: dxy, dxz, and dyz eg: dx2-y2 and dz2But the two orbitals in the eg set are now lower in energy than the three orbitals inthe t2g set, as shown in the figure below. 10. To understand the splitting of d orbitals in a tetrahedral crystal field, imagine four ligandslying at alternating corners of a cube to form a tetrahedral geometry, as shown in thefigure below. The dx2-y2 and dz2 orbitals on the metal