• Primary Valency: This refers to the oxidation state of the central metal atom in a

coordination complex. It is typically satisfied by anions and is non-directional.

• Secondary Valency: This is the number of ligands attached to the metal atom, also known
as the coordination number. It determines the geometry of the complex and is directional.

• Ligand: Ligands are species that donate electrons to metals. They act as Lewis bases,
forming coordinate covalent bonds with the central metal atom (a Lewis acid).

• Chelate: A chelate is formed when a single ligand binds to a metal ion at two or more
points, typically forming a ring structure.

• Chelating Effect: This refers to the enhanced stability of a chelate complex compared to a
complex formed by monodentate ligands (ligands that bind at only one point). When a
chelating ligand forms a complex, it results in a more stable complex.

Regarding the structure of EDTA: The notes state "EDTA -> Mn2- 6 bonds", which correctly
implies that EDTA is a hexadentate ligand, meaning it forms six bonds to the metal ion. While
your diagram in attempts to show parts of EDTA (OOC-CH2, H2N-CH2), it's incomplete and
not the full structure of Ethylenediaminetetraacetic acid. EDTA typically binds through two
nitrogen atoms and four oxygen atoms from the carboxylate groups.

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Isomerism in Coordination Complexes

Isomerism is a crucial aspect of coordination chemistry, and your notes accurately identify
the main categories.

1. Structural Isomerism These isomers have the same chemical formula but different
connectivity or arrangement of atoms. The sources list several types:

• Ionisation Isomerism: These isomers produce different ions when dissolved in water, even
though their overall composition is the same. This occurs when a counter-ion in the formula
can also act as a ligand, and vice versa.

◦ Example: [Cr(H2O)5SO4]Br and [Cr(NH3)5Br]SO4. In the first, bromide is the counter-ion;

• Hydrate Isomerism: A specific type of ionisation isomerism where water molecules are
exchanged between the coordination sphere and the lattice (as water of crystallization).

◦ Example: Violet [CrCl3(H2O)6] and green [Cr(H2O)5Cl]Cl2.H2O. The first has all six water
molecules coordinated; the second has five coordinated waters and one water of
crystallization, with one chloride ion moving into the coordination sphere.
• Linkage Isomerism: Occurs with ambidentate ligands, which can bind to the metal through
different donor atoms. For instance, the nitrite ion (NO2-) can bind via nitrogen (nitro-) or
oxygen (nitrito-).

• Coordination Isomerism: Occurs in complexes where both the cation and anion are
complex ions, and the ligands are exchanged between the two complex ions.

◦ Example: [Co(NH3)6]3+[Cr(CN)6]3- and [Cr(NH3)6]3+[Co(CN)6]3-.

2. Stereoisomerism These isomers have the same connectivity but different spatial
arrangements of atoms. The sources identify two types:

• Geometric Isomerism: Arises from different spatial arrangements of ligands around the
central metal ion.

◦ Cis and Trans isomers: Common in square planar and octahedral complexes. In cis
isomers, identical ligands are adjacent; in trans, they are opposite.

◦ Facial (fac) and Meridional (mer) isomers: Specific to octahedral complexes of the type
MA3B3. A fac isomer has three identical ligands on one face of the octahedron (forming a
triangle), while a mer isomer has the three identical ligands in a plane that includes the
central metal ion (along a meridian). Diagrams are provided for both.

• Optical Isomerism: Occurs when a molecule is non-superimposable on its mirror image,

◦ Chiral compounds are those where, if there's a carbon atom, all four groups attached to
it are different. The diagrams in illustrate non-superimposable mirror images.

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Crystal Field Theory (CFT)

CFT is a model used to explain the bonding, color, and magnetic properties of transition
metal complexes.

• Crystal Field Splitting in Octahedral Complexes:

◦ In an octahedral complex, the five degenerate d-orbitals of the central metal ion split
into two energy levels when ligands approach.

◦ The t2g set (dxy, dyz, dxz) is lowered in energy, while the eg set (dx2-y2, dz2) is raised in
energy. This splitting occurs because of the electrostatic repulsion between the d-electrons
of the metal and the lone pairs of electrons from the ligands.

◦ The energy difference between the t2g and eg sets is called the crystal field splitting
energy, denoted as Δo (for octahedral complexes).
 ◦ The t2g orbitals are stabilized by -0.4Δo relative to the barycenter (average energy), and
the eg orbitals are destabilized by +0.6Δo.

◦ This energy splitting is dependent on the strength of the ligand, which in turn influences
the complex's color and magnetic properties.

• Electron Filling and Spin States:

◦ Electrons will preferentially fill the lower energy t2g orbitals first.

◦ The distribution of electrons in the t2g and eg orbitals determines whether a complex is
high spin or low spin. This is dictated by the relative magnitudes of Δo and the pairing
energy (P).

▪ Pairing Energy (P): The energy required to place two electrons with opposite spins into
the same orbital.

▪ If Δo > P (strong field ligands), electrons will preferentially pair up in the lower energy
t2g orbitals before occupying the higher energy eg orbitals. This leads to low spin
complexes, which have fewer unpaired electrons (or none) and often exhibit more vibrant
colors.

▪ If Δo < P (weak field ligands), electrons will first singly occupy all available t2g orbitals,
then move to singly occupy eg orbitals, and only then will pairing occur. This results in high
spin complexes, which have more unpaired electrons and generally larger magnetic
moments. Weak field ligands can form complexes that are less intensely colored or
sometimes described as "colorless" compared to strong field analogues, though this is a
generalization.

• Octahedral and Tetrahedral Energy Diagrams:

◦ Octahedral: t2g (lower energy, -0.4Δo) and eg (higher energy, +0.6Δo).

◦ Tetrahedral: The splitting pattern is inverted and smaller than in octahedral complexes.
The e set (dx2-y2, dz2) is lowered in energy, and the t2 set (dxy, dyz, dxz) is raised in energy.

◦ For tetrahedral complexes, the energy difference is Δt. The e orbitals are stabilized by -
0.6Δt and the t2 orbitals are destabilized by +0.4Δt.

◦ A key relationship is Δt = (4/9)Δo. This means tetrahedral splitting is significantly smaller

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Spectrochemical Series and its Importance

The spectrochemical series is an experimentally determined list that ranks ligands based on
their ability to cause crystal field splitting (i.e., their strength).
• Order (Weak Field to Strong Field): I- < Br- < F- < H2O < NH3 < en < PPh3 < CN- < CO.

• Weak Field Ligands (e.g., I-, Br-, F-, H2O) cause less splitting (small Δo) and generally favor
high spin complexes.

• Strong Field Ligands (e.g., NH3, en, CN-, CO) cause more splitting (large Δo) and generally
favor low spin complexes.

Importance: The spectrochemical series is crucial for predicting:

• Color of a complex (larger splitting means absorption of higher energy light, affecting
observed color).

• Stability of a complex.

• Magnetic properties (high spin vs. low spin).

• It also guides the choice of ligands in inorganic chemistry for synthesis and reactivity.

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Crystal Field Stabilization Energy (CFSE) Calculation

CFSE is the stabilization energy gained by a metal ion when it forms a complex, due to the
preferential occupation of lower energy d-orbitals.

• Octahedral CFSE Formula: CFSE = (-0.4 n_t2g + 0.6 n_eg) Δo.

◦ Where n_t2g is the number of electrons in t2g orbitals, and n_eg is the number of
electrons in eg orbitals.

• Tetrahedral CFSE Formula: CFSE = (-0.6 n_e + 0.4 n_t2) Δt.

◦ Where n_e is the number of electrons in e orbitals, and n_t2 is the number of electrons
in t2 orbitals.

Examples from your notes:

• For [Fe(CN)6]4-: Fe is in +2 oxidation state (d6). CN- is a strong field ligand, so it's a low
spin complex. All 6 electrons are paired in the t2g orbitals (t2g^6 eg^0).

◦ CFSE = (-0.4 * 6 + 0.6 * 0) Δo = -2.4 Δo.

• For [Cr(NH3)6]2+: Cr is in +2 oxidation state (d4). NH3 is typically a strong field ligand for
2+ ions. However, the calculation provided seems to assume high spin (t2g^3 eg^1).

◦ If high spin: CFSE = (-0.4 * 3 + 0.6 * 1) Δo = (-1.2 + 0.6) Δo = -0.6 Δo.

• For [Fe(CN)6]3-: Fe is in +3 oxidation state (d5). CN- is a strong field ligand, so it's a low
spin complex. All 5 electrons are paired in t2g orbitals (t2g^5 eg^0).

◦ CFSE = (-0.4 * 5 + 0.6 * 0) Δo = -2.0 Δo.

◦ CFSE = (-0.4 * 3 + 0.6 * 2) Δo = (-1.2 + 1.2) Δo = 0 Δo.

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Magnetic Properties

The magnetic properties of coordination complexes are determined by the presence and
number of unpaired electrons.

• Magnetic Moment (Spin-Only Formula):

◦ The magnetic moment (M) is calculated using the formula: M = √[n(n+2)].

◦ Where n = number of unpaired electrons. The unit for magnetic moment is Bohr
Magnetons (BM).

◦ High spin complexes tend to have a larger magnetic moment due to a greater number
of unpaired electrons.

Types of Magnetism:

• Paramagnetism: Occurs in substances with unpaired electrons. These substances are

◦ Examples: [CrCl6]3- (n=3), [Ni(H2O)6]2+ (n=2), [FeF6]3- (n=5), [CoF6]3- (n=4).

• Diamagnetism: Occurs in substances where all electrons are paired. These substances are
weakly repelled by an external magnetic field.

◦ Examples: [Co(CN)6]3- (n=0), [Fe(CN)6]4- (n=0).

• Your notes also mention Ferromagnetic (e.g., Fe oxide) and Antiferromagnetic (e.g., MnO)
materials. These refer to bulk magnetic properties where there is a strong cooperative
alignment or anti-alignment of magnetic moments, respectively, in the solid state.

Magnetic Moment Calculations from your notes:

• [CrCl6]3-: Cr is +3 (d3). In an octahedral field, d3 is always high spin (t2g^3 eg^0), so n = 3

◦ M = √[3(3+2)] = √15 ≈ 3.87 BM.

• [Co(CN)6]3-: Co is +3 (d6). CN- is a strong field ligand, so it's low spin (t2g^6 eg^0). n = 0
unpaired electrons.

◦ M = √[0(0+2)] = 0 BM (Diamagnetic).
• [Ni(H2O)6]2+: Ni is +2 (d8). In an octahedral field, d8 is always high spin (t2g^6 eg^2). n = 2
unpaired electrons.

◦ M = √[2(2+2)] = √8 ≈ 2.83 BM.

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Charge Transfer Transitions (LMCT / MLCT)

Charge transfer (CT) transitions involve the transfer of an electron from predominantly
ligand-based orbitals to predominantly metal-based orbitals (LMCT) or vice versa (MLCT).
These transitions are often very intense and contribute significantly to the color of
complexes.

• Ligand-to-Metal Charge Transfer (LMCT):

◦ This occurs when an electron is transferred from a ligand-based orbital to a metal-based

◦ Conditions for LMCT:

1. High oxidation state of the metal. A metal with a higher positive charge has a
stronger pull on electrons and can more easily accept an electron.

2. Ligands with easily oxidizable electrons (i.e., good electron donors or low
electronegativity). The note states "Low electron affinity (ligand)", which is slightly
ambiguous. It is generally the metal that has a high electron affinity (or a low-lying empty
orbital) for LMCT, and the ligand that has high-lying occupied orbitals (is easily oxidized).

• Metal-to-Ligand Charge Transfer (MLCT):

◦ This occurs when an electron is transferred from a metal-based orbital to a ligand-based

◦ Conditions for MLCT:

1. Low oxidation state of the metal. A metal in a lower oxidation state can more easily
donate an electron.

2. Ligands with low-lying empty orbitals (e.g., π* antibonding orbitals), allowing them
to accept an electron (i.e., good electron acceptors or high electron affinity). Your note
correctly identifies "High electron affinity (ligand)".

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Color of Complexes

The color of coordination complexes arises primarily from d-d transitions (electron moving
between t2g and eg orbitals) and charge transfer transitions.
• When a complex absorbs light, an electron is promoted to a higher energy level. The color
observed is the complementary color to the light absorbed.

• For example, if a complex absorbs yellow light, it will appear violet. The color wheel in your
notes illustrates this concept.

• Strong field ligands cause a larger Δo, leading to the absorption of higher energy (shorter
wavelength) light, often resulting in more vibrant colors.

• Weak field ligands cause a smaller Δo, absorbing lower energy (longer wavelength) light.

Example Calculation:

• For a complex like [Ni(H2O)6]2+, if it absorbs light with a wavenumber of 20000 cm-1:

◦ Wavelength (λ) = 1 / (Wavenumber) = 1 / 20000 cm-1 = 0.00005 cm = 500 nm.

◦ This corresponds to the absorption of green-blue light.

• The energy of this absorbed light can be calculated using E = hc/λ. Your note shows a
conversion to 2.48 eV, which is a typical energy value for visible light absorption.