Transition Metals Crash Course

Hey there, so this is the periodic table.


Most of the elements you see here are metals. What are metals like? If you recall from my previous post, Chemical Bonds, Especially Covalent Ones, you will remember a bit of metallic bonding.

Electrons are loosely held and they move through a lattice, lending some physical characteristics to the solid, such as high luster (ability to absorb and reemit all wavelengths of light), good conductors of electricity and heat (due to delocalized electrons), malleability, ductility, and high density (close packed arrangement).

Chemically, metals form positive ions and have positive oxidation states, they tend to form ionic compounds, and metals form oxides which produce basic metal hydroxides when reacted with water.

Today, I’m excited to announce that we will be focusing on transition metals! They are the huge yellow block of metals in the periodic table above from group 3 to 11.


So what are the transition metals? Basically, they all have an incompletely filled d sub-shell, or can give rise to cations with incomplete d sub-shells.

Even though Scandium (Sc) appears to be a transition metal, it is not, because the ion of Sc is 3+ and lacks electrons.

Some characteristics of transition metals:

  1. Small, compact atoms
  2. Paramagnetism (due to unpaired d orbitals)
  3. Variable oxidation states (Mn has +2,+3,+4,+6,+6,+7, for example)
  4. Colored compounds
  5. Stable complexes
  6. Catalytic activity

Complex Ions and Coordination Compounds

A complex is a molecule or ion that consists of a central metal atom or ion surrounded by a number of other atoms or molecules called ligands, which donate electrons to the central atom or ion.

Usually the central atom/ion is a Lewis acid (electron deficient, i.e. transition metal ions), which are surrounded by Lewis bases (electron donors, for example NH3, H2O, Cl-, CN-).

Neutral compounds that contain complexes are called coordination compounds.

The metal-ligand bond in a complex is considered to be a coordinate covalent bond (i.e. dative bond).The ligand’s coordinating atom orients its lone pair towards the central metal atom, and the coordination number is the number of these ligand atoms in contact with the central metal atom.

Like in VSEPR theory, the geometrical structure for metal complexes depends on the coordination number (two is linear, four is tetrahedral, six is octahedral, etc.)

A ligand is mondentate, bidentate, or tridentate depending on whether it has one, two, or three corrdinating atoms (it can act like a claw). A polydentate ligand that binds a metal atom in two or more positions is called a chelating agent.

Nomenclature (Naming Complexes)

You can safely ignore this part unless you are taking a course in which you need to know how to name coordination compounds using IUPAC standards.

  1. The cation is named before the anion.
  2. Within a complex ion, the ligands are named first in alphabetical order, and the metal ion is named last.
    1. Anionic ligands are given names ending in -o: Cl- (chloro), CN- (cyano), OH-, (hydroxo), SO4 2- (sulfato), NO2- (nitro), NO3 (nitrato)
    2. Neutral ligands keep their own names.
    3. A few ligands have special names: H2O (aqua), NH3 (ammine), CO (carbonyl), NO (nitrosyl)
    4. Greek prefixes di, tri, and tetra indicate the number of identical ligands present.
    5. The prefixes bis, tris, and tetrakis are used instead of the Greek prefixes to avoid confusion.
    6. Prefixes are disregarded in alphabetizing the ligands.
  3. The name of the central atom is followed by its oxidation state designated by a Roman numeral enclosed in parentheses. In a positive or neutral complex the metal is given its usual name. If the complex ion has a negative charge, the name of the central atom ends in ate, and the Latin root is used for metals whose symbols have a Latin derivation (for example, ferrate for iron, aluminate for aluminum).
  4. The name of the complex is written without spaces. If the complex is charged, the name ends with the word ion.

Isomerism in Metal Complexes

Isomers are compounds with the same chemical composition but different arrangements of atoms. There are two types: (1) structural and (2) stereoisomers.

Structural Isomers

The atoms are joined in a different sequence, so each structural isomer has a different set of bonds.


These have the same bonds but they differ in the spatial arrangement of these bonds. There are two types: (1) geometric isomers and (2) enantiomers (optical isomers).

  1. Geometric Isomers
    These are called cis (similar ligands on same side) and trans (similar ligands on opposite sides) isomers.

  2. Enantiomers
    These are stereoisomers that are chiral, or nonsuperimposable mirror images. They are also known as optical isomers because the dextrorotary (d) isomer rotates plane-polarized light to the right, and the levoratatory (l) isomer rotates it to the left. In an equimolar mixture of two enantiomers the net rotation is zero and the mixture is racemic.

Theories of Bonding in Metal Complexes

  1. Crystal Field Theory is a simplified theory that approximates ligands as point charges, which we will be covering today.
  2. Ligand Field Theory is an extension of molecular orbital theory to cover transition metal complexes, which is beyond the scope of this course.

These theories seek to explain (1) absorption spectra (color) (2) magnetic properties (paramagnetism) and (3) geometrical arrangement of atoms.

Crystal Field Theory

In this theory we assume that the bonding between the central metal ion and its ligands is purely electrostatic and ionic in nature. We also assume that the ligand electrons remain on the ligands and the metral electrons remain on the metal.

Like we mentioned above, the lone pairs on the ligands are treated as point negative charges.


(1) Octahedral Complexes
A transition metal ion in free space would have five d orbitals with the same energy.
Now imagine six ligands placed symmetrically around this central ion.
Bring the ligands closer to the central ion. The energies of all the d orbitals will rise due to electrostatic repulsion.
The electrons in the two eg orbitals will feel an even stronger repulsion than the other three orbitals (t2g) because these orbitals are concentrated along the coordinate axes where the ligands are situated.
This will cause these two eg orbitals to lie higher in energy than the three t2g orbitals.
This difference in energy is called the ligand field splitting or crystal field splitting energy (Δo)
In octahedral complexes, the energy of the three t2g orbitals is lowered by 2/5 Δo and the energy of the two eg orbitals is raised by 3/5 Δo, relative the the energy of the d orbitals in the presence of a uniform spherical charge distribution.

(2) Tetrahedral Complexes
The crystal field splitting for tetrahedral complexes is reverse that of the octahedral splitting because the three t2 orbitals point more directly at the ligands than the two e orbitals.

Absorption Spectra

Transition metal complexes are very colorful!

This can be attributed to the complexes absorbing certain frequencies in the visible spectrum. The absorption of light is accompanied by the excitation of an electron in one of the filled lower energy t2g orbitals to a higher energy empty eg orbital.

These electronic transitions are called d-d transitions. The energy of the color most strongly absorbed is equal to Δo, so

Δo = hv = hc / λ

Note that the wavelength of the color absorbed is not the color we observe. The observed color of the complex is the complementary color to the color absorbed by the complex.

Spectrochemical Series

A ligand that produces a large splitting energy is a strong-field ligand. One that produces a small splitting is called a weak-field ligand.

A list of ligands in order of field strength is called a spectrochemical series:

I- < Br- < Cl- < F- < OH- < C2O4 2- < H2O < NCS- < NH3 < ethylenediamine < NO2 2- < CN- and CO

Magnetic Properties

If the electron configuration is 3d1, 3d2, or 3d3, there is no ambiguity in the distribution of electrons because they all go into the t2g orbitals.

If the electron configuration is 3d4, 3d5, 3d6, or 3d7, the electrons can go in either the t2g or eg orbitals.

If Δo is large (strong-field ligand), pair formation in the lower t2g orbitals occurs (low-spin complex).

If Δo is small (weak-field ligand), unpaired electrons end up in the higher eg orbitals (high-spin complex) because of the lower electron-electron repulsion associated with unpaired electrons.

These unpaired electrons make a substance paramagnetic, or weakly attracted by a magnetic field.

That’s all I have for you folks! Check back in a few days’ time to get an overview of chemical kinetics. Bye! 🙂

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