1. Ligand Field Strength
2. Is Dmg A Strong Field Ligand In Chemistry

Sep 05, 2019  Thanks for A2A!!! You should learn the spectrochemical series to know which are weak field ligands and which are strong field ligands. Weak field ligands: I-, Br-, SCN-, Cl-, F-, OH-, NO2-, H2O. These ligands doesn’t help in the pairing of. Start studying Strong and Weak Field Ligands. Learn vocabulary, terms, and more with flashcards, games, and other study tools. Dmg is a strong field ligand or a weak field ligand? Share with your friends.

A spectrochemical series is a list of ligands ordered on ligand strength and a list of metal ions based on oxidation number, group and its identity. In crystal field theory, ligands modify the difference in energy between the d orbitals (Δ) called the ligand-field splitting parameter for ligands or the crystal-field splitting parameter, which is mainly reflected in differences in color of similar metal-ligand complexes.

Spectrochemical series of ligands[edit]

The spectrochemical series was first proposed in 1938 based on the results of absorption spectra of cobalt complexes.^[1]

A partial spectrochemical seriesWhat are dmg files should i save. listing of ligands from small Δ to large Δ is given below. (For a table, see the ligand page.)

I⁻ < Br⁻ < S²⁻ < SCN⁻ (S–bonded) < Cl⁻ < N₃⁻ < F⁻< NCO⁻ < OH⁻ < C₂O₄²⁻ < O²⁻< H₂O < acac⁻ (acetylacetonate) < NCS⁻ (N–bonded) < CH₃CN < gly (glycine) < py (pyridine) < NH₃ < en (ethylenediamine) < bipy (2,2'-bipyridine) < phen (1,10-phenanthroline) < NO₂⁻ < PPh₃ < CN⁻ < CO

Weak field ligand: H₂O,F-,Cl-,OH-Strong field ligand: CO,CN-,NH₃,PPh₃

Ligands arranged on the left end of this spectrochemical series are generally regarded as weaker ligands and cannot cause forcible pairing of electrons within the 3d level, and thus form outer orbital octahedral complexes that are high spin. On the other hand, ligands lying at the right end are stronger ligands and form inner orbital octahedral complexes after forcible pairing of electrons within 3d level and hence are called low spin ligands.

However, keep in mind that 'the spectrochemical series is essentially backwards from what it should be for a reasonable prediction based on the assumptions of crystal field theory.'^[2] This deviation from crystal field theory highlights the weakness of crystal field theory's assumption of purely ionic bonds between metal and ligand.

The order of the spectrochemical series can be derived from the understanding that ligands are frequently classified by their donor or acceptor abilities. Some, like NH₃, are σ bond donors only, with no orbitals of appropriate symmetry for π bonding interactions. Bonding by these ligands to metals is relatively simple, using only the σ bonds to create relatively weak interactions. Another example of a σ bonding ligand would be ethylenediamine, however ethylenediamine has a stronger effect than ammonia, generating a larger ligand field split, Δ.

Ligands that have occupied p orbitals are potentially π donors. These types of ligands tend to donate these electrons to the metal along with the σ bonding electrons, exhibiting stronger metal-ligand interactions and an effective decrease of Δ. Most halide ligands as well as OH⁻ are primary examples of π donor ligands.

When ligands have vacant π* and d orbitals of suitable energy, there is the possibility of pi backbonding, and the ligands may be π acceptors. This addition to the bonding scheme increases Δ. Ligands that do this very effectively include CN⁻, CO, and many others.^[3]

Spectrochemical series of metals[edit]

The metal ions can also be arranged in order of increasing Δ, and this order is largely independent of the identity of the ligand.^[4]

Mn²⁺ < Ni²⁺ < Co²⁺ < Fe²⁺ < V²⁺ < Fe³⁺ < Cr³⁺ < V³⁺ < Co³⁺

In general, it is not possible to say whether a given ligand will exert a strong field or a weak field on a given metal ion. However, when we consider the metal ion, the following two useful trends are observed:

• Δ increases with increasing oxidation number, and
• Δ increases down a group.^[4]

See also[edit]

References[edit]

• Zumdahl, Steven S. Chemical Principles Fifth Edition. Boston: Houghton Mifflin Company, 2005. Pages 550-551 and 957-964.
• D. F. Shriver and P. W. Atkins Inorganic Chemistry 3rd edition, Oxford University Press, 2001. Pages: 227-236.
• James E. Huheey, Ellen A. Keiter, and Richard L. Keiter Inorganic Chemistry: Principles of Structure and Reactivity 4th edition, HarperCollins College Publishers, 1993. Pages 405-408.

1. ^R. Tsuchida (1938). 'Absorption Spectra of Co-ordination Compounds. I.'Bull. Chem. Soc. Jpn. 13 (5). doi:10.1246/bcsj.13.388.
2. ^7th page of http://science.marshall.edu/castella/chm448/chap11.pdf
3. ^Miessler, Gary; Tarr, Donald (2011). Inorganic Chemistry (4th ed.). Prentice Hall. pp. 395–396. ISBN978-0-13-612866-3.
4. ^ ^abhttp://www.everyscience.com/Chemistry/Inorganic/Crystal_and_Ligand_Field_Theories/b.1013.php

Ligand field theory (LFT) describes the bonding, orbital arrangement, and other characteristics of coordination complexes.^[1][2][3] It represents an application of molecular orbital theory to transition metal complexes. A transition metal ion has nine valence atomic orbitals - consisting of five nd, three (n+1)p, and one (n+1)s orbitals. These orbitals are of appropriate energy to form bonding interaction with ligands. The LFT analysis is highly dependent on the geometry of the complex, but most explanations begin by describing octahedral complexes, where six ligands coordinate to the metal. Other complexes can be described by reference to crystal field theory.^[4]

History[edit]

Ligand field theory resulted from combining the principles laid out in molecular orbital theory and crystal field theory, which describes the loss of degeneracy of metal d orbitals in transition metal complexes. John Stanley Griffith and Leslie Orgel^[5] championed ligand field theory as a more accurate description of such complexes, although the theory originated in the 1930s with the work on magnetism of John Hasbrouck Van Vleck. Griffith and Orgel used the electrostatic principles established in crystal field theory to describe transition metal ions in solution and used molecular orbital theory to explain the differences in metal-ligand interactions, thereby explaining such observations as crystal field stabilization and visible spectra of transition metal complexes. In their paper, they proposed that the chief cause of color differences in transition metal complexes in solution is the incomplete d orbital subshells.^[5] That is, the unoccupied d orbitals of transition metals participate in bonding, which influences the colors they absorb in solution. In ligand field theory, the various d orbitals are affected differently when surrounded by a field of neighboring ligands and are raised or lowered in energy based on the strength of their interaction with the ligands.^[5]

Bonding[edit]

σ-bonding (sigma bonding)[edit]

In an octahedral complex, the molecular orbitals created by coordination can be seen as resulting from the donation of two electrons by each of six σ-donor ligands to the d-orbitals on the metal. In octahedral complexes, ligands approach along the x-, y- and z-axes, so their σ-symmetry orbitals form bonding and anti-bonding combinations with the d_z² and d_x²−y² orbitals. The d_xy, d_xz and d_yz orbitals remain non-bonding orbitals. Some weak bonding (and anti-bonding) interactions with the s and p orbitals of the metal also occur, to make a total of 6 bonding (and 6 anti-bonding) molecular orbitals

In molecular symmetry terms, the six lone-pair orbitals from the ligands (one from each ligand) form six symmetry adapted linear combinations (SALCs) of orbitals, also sometimes called ligand group orbitals (LGOs). The irreducible representations that these span are a_1g, t_1u and e_g. The metal also has six valence orbitals that span these irreducible representations - the s orbital is labeled a_1g, a set of three p-orbitals is labeled t_1u, and the d_z² and d_x²−y² orbitals are labeled e_g. The six σ-bonding molecular orbitals result from the combinations of ligand SALCs with metal orbitals of the same symmetry.

π-bonding (pi bonding)[edit]

π bonding in octahedral complexes occurs in two ways: via any ligand p-orbitals that are not being used in σ bonding, and via any π or π^* molecular orbitals present on the ligand.

In the usual analysis, the p-orbitals of the metal are used for σ bonding (and have the wrong symmetry to overlap with the ligand p or π or π^* orbitals anyway), so the π interactions take place with the appropriate metal d-orbitals, i.e. d_xy, d_xz and d_yz. These are the orbitals that are non-bonding when only σ bonding takes place.

One important π bonding in coordination complexes is metal-to-ligand π bonding, also called π backbonding. It occurs when the LUMOs (lowest unoccupied molecular orbitals) of the ligand are anti-bonding π^* orbitals. These orbitals are close in energy to the d_xy, d_xz and d_yz orbitals, with which they combine to form bonding orbitals (i.e. orbitals of lower energy than the aforementioned set of d-orbitals). The corresponding anti-bonding orbitals are higher in energy than the anti-bonding orbitals from σ bonding so, after the new π bonding orbitals are filled with electrons from the metal d-orbitals, Δ_O has increased and the bond between the ligand and the metal strengthens. The ligands end up with electrons in their π^* molecular orbital, so the corresponding π bond within the ligand weakens.

The other form of coordination π bonding is ligand-to-metal bonding. This situation arises when the π-symmetry p or π orbitals on the ligands are filled. They combine with the d_xy, d_xz and d_yz orbitals on the metal and donate electrons to the resulting π-symmetry bonding orbital between them and the metal. The metal-ligand bond is somewhat strengthened by this interaction, but the complementary anti-bonding molecular orbital from ligand-to-metal bonding is not higher in energy than the anti-bonding molecular orbital from the σ bonding. It is filled with electrons from the metal d-orbitals, however, becoming the HOMO (highest occupied molecular orbital) of the complex. For that reason, Δ_O decreases when ligand-to-metal bonding occurs.

The greater stabilization that results from metal-to-ligand bonding is caused by the donation of negative charge away from the metal ion, towards the ligands. This allows the metal to accept the σ bonds more easily. The combination of ligand-to-metal σ-bonding and metal-to-ligandπ-bonding is a synergic effect, as each enhances the other. Potion of diminution rare dmg 187.

As each of the six ligands has two orbitals of π-symmetry, there are twelve in total. The symmetry adapted linear combinations of these fall into four triply degenerate irreducible representations, one of which is of t_2g symmetry. The d_xy, d_xz and d_yz orbitals on the metal also have this symmetry, and so the π-bonds formed between a central metal and six ligands also have it (as these π-bonds are just formed by the overlap of two sets of orbitals with t_2g symmetry.)

Role of metal p-orbitals[edit]

Current computational findings suggest valence p orbitals on the metal participate in metal-ligand bonding, albeit weakly.^[6] Some new theoretical treatments do not count the metal p-orbitals in metal-ligand bonding,^[7] although these orbitals are still included as polarization functions. This model has yet to be adopted by the general chemistry community.

High and low spin and the spectrochemical series[edit]

The six bonding molecular orbitals that are formed are 'filled' with the electrons from the ligands, and electrons from the d-orbitals of the metal ion occupy the non-bonding and, in some cases, anti-bonding MOs. The energy difference between the latter two types of MOs is called Δ_O (O stands for octahedral) and is determined by the nature of the π-interaction between the ligand orbitals with the d-orbitals on the central atom. As described above, π-donor ligands lead to a small Δ_O and are called weak- or low-field ligands, whereas π-acceptor ligands lead to a large value of Δ_O and are called strong- or high-field ligands. Ligands that are neither π-donor nor π-acceptor give a value of Δ_O somewhere in-between.

The size of Δ_O determines the electronic structure of the d⁴ - d⁷ ions. In complexes of metals with these d-electron configurations, the non-bonding and anti-bonding molecular orbitals can be filled in two ways: one in which as many electrons as possible are put in the non-bonding orbitals before filling the anti-bonding orbitals, and one in which as many unpaired electrons as possible are put in. The former case is called low-spin, while the latter is called high-spin. A small Δ_O can be overcome by the energetic gain from not pairing the electrons, leading to high-spin. When Δ_O is large, however, the spin-pairing energy becomes negligible by comparison and a low-spin state arises.

Ligand Field Strength

The spectrochemical series is an empirically-derived list of ligands ordered by the size of the splitting Δ that they produce. It can be seen that the low-field ligands are all π-donors (such as I⁻), the high field ligands are π-acceptors (such as CN⁻ and CO), and ligands such as H₂O and NH₃, which are neither, are in the middle.

I⁻ < Br⁻ < S²⁻ < SCN⁻ < Cl⁻ < NO₃⁻ < N₃⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < CH₃CN < py (pyridine) < NH₃ < en (ethylenediamine) < bipy (2,2'-bipyridine) < phen (1,10-phenanthroline) < NO₂⁻ < PPh₃ < CN⁻ < CO

See also[edit]

References[edit]

1. ^Ballhausen, Carl Johan,'Introduction to Ligand Field Theory',McGraw-Hill Book Co., New York, 1962
2. ^Griffith, J.S. (2009). The Theory of Transition-Metal Ions (re-issue ed.). Cambridge University Press. ISBN978-0521115995.
3. ^Schläfer, H. L.; Gliemann, G. 'Basic Principles of Ligand Field Theory' Wiley Interscience: New York; 1969
4. ^G. L. Miessler and D. A. Tarr “Inorganic Chemistry” 3rd Ed, Pearson/Prentice Hall publisher, ISBN0-13-035471-6.
5. ^ ^abcGriffith, J.S. and L.E. Orgel. 'Ligand Field Theory'.Q. Rev. Chem. Soc. 1957, 11, 381-393
6. ^Frenking, Gernot; Shaik, Sason, eds. (May 2014). 'Chapter 7: Chemical bonding in Transition Metal Compounds'. The Chemical Bond: Chemical Bonding Across the Periodic Table. Wiley -VCH. ISBN978-3-527-33315-8.
7. ^C. R. Landis, F. Weinhold (2007). 'Valence and extra-valence orbitals in main group and transition metal bonding'. Journal of Computational Chemistry. 28 (1): 198–203. doi:10.1002/jcc.20492. PMID17063478.

External links[edit]

• Crystal-field Theory, Tight-binding Method, and Jahn-Teller Effect in E. Pavarini, E. Koch, F. Anders, and M. Jarrell (eds.): Correlated Electrons: From Models to Materials, Jülich 2012, ISBN978-3-89336-796-2

Is Dmg A Strong Field Ligand In Chemistry