Spin states when describing transition metalcoordination complexes refers to the potential spin configurations of the central metal's d electrons. For several oxidation states, metals can adopt high-spin and low-spin configurations. The ambiguity only applies to first row metals, because second- and third-row metals are invariably low-spin. These configurations can be understood through the two major models used to describe coordination complexes; crystal field theory and ligand field theory (a more advanced version based on molecular orbital theory).[1]
The Δ splitting of the d orbitals plays an important role in the electron spin state of a coordination complex. Three factors affect Δ: the period (row in periodic table) of the metal ion, the charge of the metal ion, and the field strength of the complex's ligands as described by the spectrochemical series. Only octahedral complexes of first row transition metals adopt high-spin states.
In order for low spin splitting to occur, the energy cost of placing an electron into an already singly occupied orbital must be less than the cost of placing the additional electron into an eg orbital at an energy cost of Δ. If the energy required to pair two electrons is greater than the energy cost of placing an electron in an eg, Δ, high spin splitting occurs.
If the separation between the orbitals is large, then the lower energy orbitals are completely filled before population of the higher orbitals according to the Aufbau principle. Complexes such as this are called "low-spin" since filling an orbital matches electrons and reduces the total electron spin. If the separation between the orbitals is small enough then it is easier to put electrons into the higher energy orbitals than it is to put two into the same low-energy orbital, because of the repulsion resulting from matching two electrons in the same orbital. So, one electron is put into each of the five d orbitals before any pairing occurs in accord with Hund's rule resulting in what is known as a "high-spin" complex. Complexes such as this are called "high-spin" since populating the upper orbital avoids matches between electrons with opposite spin.
High-spin [FeBr6]3− crystal field diagram
The charge of the metal center plays a role in the ligand field and the Δ splitting. The higher the oxidation state of the metal, the stronger the ligand field that is created. In the event that there are two metals with the same d electron configuration, the one with the higher oxidation state is more likely to be low spin than the one with the lower oxidation state; for example, Fe2+ and Co3+ are both d6; however, the higher charge of Co3+ creates a stronger ligand field than Fe2+. All other things being equal, Fe2+ is more likely to be high spin than Co3+.
Ligands also affect the magnitude of Δ splitting of the d orbitals according to their field strength as described by the spectrochemical series. Strong-field ligands, such as CN− and CO, increase the Δ splitting and are more likely to be low-spin. Weak-field ligands, such as I− and Br− cause a smaller Δ splitting and are more likely to be high-spin.
Some octahedral complexes exhibit spin crossover, where the high and low spin states exist is dynamic equilibrium.
Light-induced spin-crossover of [Fe(pyCH2NH2)3]2+, which switches from high and low-spin.[2]
Tetrahedral complexesEdit
Fe(4-norbornyl)4 is a rare example of a low-spin tetrahedral complex.
The Δ splitting energy for tetrahedral metal complexes (four ligands), Δtet is smaller than that for an octahedral complex. Consequently, tetrahedral complexes are almost always high spin[3] Examples of low spin tetrahedral complexes include Fe(2-norbornyl)4,[4] [Co(4-norbornyl)4]+, and the nitrosyl complex Cr(NO)((N(tms)2)3.
Square planar complexesEdit
Many d8 complexes of the first row metals exist in tetrahedral or square planar geometry. In some cases these geometries exist in measurable equilibria. For example, dichlorobis(triphenylphosphine)nickel(II) has been crystallized in both tetrahedral and square planar geometries.[5]
Ligand field theory vs crystal field theoryEdit
In terms of d-orbital splitting, ligand field theory (LFT) and crystal field theory (CFT) give similar results. CFT is an older, simpler model that treats ligands as point charges. LFT is more chemical, emphasizes covalent bonding and accommodates pi-bonding explicitly.
High-spin and low-spin systemsEdit
In the case of octahedral complexes, the question of high spin vs low spin first arises for d4, since it has more than the 3 electrons to fill the non-bonding d orbitals according to ligand field theory or the stabilized d orbitals according to crystal field splitting.
All complexes of second and third row metals are low-spin.
d4
Octahedral high-spin: 4 unpaired electrons, paramagnetic, substitutionally labile. Includes Cr2+ (many complexes assigned as Cr(II) are however Cr(III) with reduced ligands[6]), Mn3+.
Generally, the rates of ligand dissociation from low spin complexes are lower than dissociation rates from high spin complexes. In the case of octahedral complexes, electrons in the eg levels are anti-bonding with respect to the metal-ligand bonds. Famous "exchange inert" complexes are octahedral complexes of d3 and low-spin d6 metal ions, illustrated respectfully by Cr3+ and Co3+.[8]
ReferencesEdit
^Miessler, Gary L.; Donald A. Tarr (1998). Inorganic Chemistry (2nd ed.). Upper Saddle River, New Jersey: Pearson Education, Inc. Pearson Prentice Hall. ISBN0-13-841891-8.
^Zumdahl, Steven (2009). "19.6 Transition Metals and Coordination Chemistry: The Crystal Field Model". Chemical Principles. Cengage Learning, Inc. ISBN978-0538734561.
^Bower, Barton K.; Tennent, Howard G. (1972). "Transition Metal Bicyclo[2.2.1]hept-1-yls". Journal of the American Chemical Society. 94 (7): 2512–2514. doi:10.1021/ja00762a056.
^Batsanov, Andrei S.; Howard, Judith A. K. (2001). "trans-Dichlorobis(triphenylphosphine)nickel(II) Bis(dichloromethane) Solvate: Redetermination at 120 K". Acta Crystallogr E. 57: 308–309. doi:10.1107/S1600536801008741. S2CID 97381117.
^Scarborough, Christopher C.; Sproules, Stephen; Doonan, Christian J.; Hagen, Karl S.; Weyhermüller, Thomas; Wieghardt, Karl (2012). "Scrutinizing Low-Spin Cr(II) Complexes". Inorganic Chemistry. 51 (12): 6969–6982. doi:10.1021/ic300882r. PMID 22676275.
^Shannon R.D. (1976). "Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides". Acta Crystallographica. A32 (5): 751-767. doi:10.1107/S0567739476001551.
^R. G. Wilkins (1991). Kinetics and Mechanism of Reactions of Transition Metal Complexes, 2nd Thoroughly Revised Edition. Weinheim: VCH. doi:10.1002/bbpc.19920960429. ISBN3-527-28389-7.
August 26, 2023
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Spin states when describing transition metal coordination complexes refers to the potential spin configurations of the central metal s d electrons For several oxidation states metals can adopt high spin and low spin configurations The ambiguity only applies to first row metals because second and third row metals are invariably low spin These configurations can be understood through the two major models used to describe coordination complexes crystal field theory and ligand field theory a more advanced version based on molecular orbital theory 1 Contents 1 High spin vs low spin 1 1 Octahedral complexes 1 2 Tetrahedral complexes 1 3 Square planar complexes 2 Ligand field theory vs crystal field theory 3 High spin and low spin systems 3 1 Ionic radii 4 Ligand exchange rates 5 ReferencesHigh spin vs low spin EditMain article Magnetochemistry Octahedral complexes Edit Low spin Fe NO2 6 3 crystal field diagramThe D splitting of the d orbitals plays an important role in the electron spin state of a coordination complex Three factors affect D the period row in periodic table of the metal ion the charge of the metal ion and the field strength of the complex s ligands as described by the spectrochemical series Only octahedral complexes of first row transition metals adopt high spin states In order for low spin splitting to occur the energy cost of placing an electron into an already singly occupied orbital must be less than the cost of placing the additional electron into an eg orbital at an energy cost of D If the energy required to pair two electrons is greater than the energy cost of placing an electron in an eg D high spin splitting occurs If the separation between the orbitals is large then the lower energy orbitals are completely filled before population of the higher orbitals according to the Aufbau principle Complexes such as this are called low spin since filling an orbital matches electrons and reduces the total electron spin If the separation between the orbitals is small enough then it is easier to put electrons into the higher energy orbitals than it is to put two into the same low energy orbital because of the repulsion resulting from matching two electrons in the same orbital So one electron is put into each of the five d orbitals before any pairing occurs in accord with Hund s rule resulting in what is known as a high spin complex Complexes such as this are called high spin since populating the upper orbital avoids matches between electrons with opposite spin High spin FeBr6 3 crystal field diagramThe charge of the metal center plays a role in the ligand field and the D splitting The higher the oxidation state of the metal the stronger the ligand field that is created In the event that there are two metals with the same d electron configuration the one with the higher oxidation state is more likely to be low spin than the one with the lower oxidation state for example Fe2 and Co3 are both d6 however the higher charge of Co3 creates a stronger ligand field than Fe2 All other things being equal Fe2 is more likely to be high spin than Co3 Ligands also affect the magnitude of D splitting of the d orbitals according to their field strength as described by the spectrochemical series Strong field ligands such as CN and CO increase the D splitting and are more likely to be low spin Weak field ligands such as I and Br cause a smaller D splitting and are more likely to be high spin Some octahedral complexes exhibit spin crossover where the high and low spin states exist is dynamic equilibrium Light induced spin crossover of Fe pyCH2NH2 3 2 which switches from high and low spin 2 Tetrahedral complexes Edit Fe 4 norbornyl 4 is a rare example of a low spin tetrahedral complex The D splitting energy for tetrahedral metal complexes four ligands Dtet is smaller than that for an octahedral complex Consequently tetrahedral complexes are almost always high spin 3 Examples of low spin tetrahedral complexes include Fe 2 norbornyl 4 4 Co 4 norbornyl 4 and the nitrosyl complex Cr NO N tms 2 3 Square planar complexes Edit Many d8 complexes of the first row metals exist in tetrahedral or square planar geometry In some cases these geometries exist in measurable equilibria For example dichlorobis triphenylphosphine nickel II has been crystallized in both tetrahedral and square planar geometries 5 Ligand field theory vs crystal field theory EditIn terms of d orbital splitting ligand field theory LFT and crystal field theory CFT give similar results CFT is an older simpler model that treats ligands as point charges LFT is more chemical emphasizes covalent bonding and accommodates pi bonding explicitly High spin and low spin systems EditIn the case of octahedral complexes the question of high spin vs low spin first arises for d4 since it has more than the 3 electrons to fill the non bonding d orbitals according to ligand field theory or the stabilized d orbitals according to crystal field splitting All complexes of second and third row metals are low spin d4 Octahedral high spin 4 unpaired electrons paramagnetic substitutionally labile Includes Cr2 many complexes assigned as Cr II are however Cr III with reduced ligands 6 Mn3 Octahedral low spin 2 unpaired electrons paramagnetic substitutionally inert Includes Cr2 Mn3 d5 Octahedral high spin 5 unpaired electrons paramagnetic substitutionally labile Includes Fe3 Mn2 Example Tris acetylacetonato iron III Octahedral low spin 1 unpaired electron paramagnetic substitutionally inert Includes Fe3 Example Fe CN 6 3 d6 Octahedral high spin 4 unpaired electrons paramagnetic substitutionally labile Includes Fe2 Co3 Examples Fe H2O 6 2 CoF6 3 Octahedral low spin no unpaired electrons diamagnetic substitutionally inert Includes Fe2 Co3 Ni4 Example Co NH3 6 3 d7 Octahedral high spin 3 unpaired electrons paramagnetic substitutionally labile Includes Co2 Ni3 Octahedral low spin 1 unpaired electron paramagnetic substitutionally labile Includes Co2 Ni3 Example Co NH3 6 2 d8 Octahedral high spin 2 unpaired electrons paramagnetic substitutionally labile Includes Ni2 Example Ni NH3 6 2 Tetrahedral high spin 2 unpaired electrons paramagnetic substitutionally labile Includes Ni2 Example NiCl4 2 Square planar low spin no unpaired electrons diamagnetic substitutionally inert Includes Ni2 Example Ni CN 4 2 Ionic radii Edit The spin state of the complex affects an atom s ionic radius For a given d electron count high spin complexes are larger 7 d4 Octahedral high spin Cr2 64 5 pm Octahedral low spin Mn3 58 pm d5 Octahedral high spin Fe3 the ionic radius is 64 5 pm Octahedral low spin Fe3 the ionic radius is 55 pm d6 Octahedral high spin Fe2 the ionic radius is 78 pm Co3 ionic radius 61 pm Octahedral low spin Includes Fe2 ionic radius 62 pm Co3 ionic radius 54 5 pm Ni4 ionic radius 48 pm d7 Octahedral high spin Co2 ionic radius 74 5 pm Ni3 ionic radius 60 pm Octahedral low spin Co2 ionic radius 65 pm Ni3 ionic radius 56 pm d8 Octahedral high spin Ni2 ionic radius 69 pm Square planar low spin Ni2 ionic radius 49 pm Ligand exchange rates EditGenerally the rates of ligand dissociation from low spin complexes are lower than dissociation rates from high spin complexes In the case of octahedral complexes electrons in the eg levels are anti bonding with respect to the metal ligand bonds Famous exchange inert complexes are octahedral complexes of d3 and low spin d6 metal ions illustrated respectfully by Cr3 and Co3 8 References Edit Miessler Gary L Donald A Tarr 1998 Inorganic Chemistry 2nd ed Upper Saddle River New Jersey Pearson Education Inc Pearson Prentice Hall ISBN 0 13 841891 8 Gutlich P 2001 Photoswitchable Coordination Compounds Coordination Chemistry Reviews 219 221 839 879 doi 10 1016 S0010 8545 01 00381 2 Zumdahl Steven 2009 19 6 Transition Metals and Coordination Chemistry The Crystal Field Model Chemical Principles Cengage Learning Inc ISBN 978 0538734561 Bower Barton K Tennent Howard G 1972 Transition Metal Bicyclo 2 2 1 hept 1 yls Journal of the American Chemical Society 94 7 2512 2514 doi 10 1021 ja00762a056 Batsanov Andrei S Howard Judith A K 2001 trans Dichlorobis triphenylphosphine nickel II Bis dichloromethane Solvate Redetermination at 120 K Acta Crystallogr E 57 308 309 doi 10 1107 S1600536801008741 S2CID 97381117 Scarborough Christopher C Sproules Stephen Doonan Christian J Hagen Karl S Weyhermuller Thomas Wieghardt Karl 2012 Scrutinizing Low Spin Cr II Complexes Inorganic Chemistry 51 12 6969 6982 doi 10 1021 ic300882r PMID 22676275 Shannon R D 1976 Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides Acta Crystallographica A32 5 751 767 doi 10 1107 S0567739476001551 R G Wilkins 1991 Kinetics and Mechanism of Reactions of Transition Metal Complexes 2nd Thoroughly Revised Edition Weinheim VCH doi 10 1002 bbpc 19920960429 ISBN 3 527 28389 7 Retrieved from https en wikipedia org w index php title Spin states d electrons amp 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