Transition Metal Properties Chemistry Tutorial

Key Concepts

• Transition metals are located in the middle of the Periodic Table and have an electron configuration filling the d-subshell as shown below:

  Periodic Table of the Elements
  s subshell		d subshell										p subshell
  H	He
  Li	Be											B	C	N	O	F	Ne
  Na	Mg											Al	Si	P	S	Cl	Ar
  K	Ca	Sc	Ti	V	Cr	Mn	Fe	Co	Ni	Cu	Zn	Ga	Ge	As	Se	Br	Kr
  Rb	Sr	Y	Zr	Nb	Mo	Tc	Ru	Rh	Pd	Ag	Cd	In	Sn	Sb	Te	I	Xe
  Cs	Ba	La	Hf	Ta	W	Re	Os	Ir	Pt	Au	Hg	Tl	Pb	Bi	Po	At	Rn
  		Transition Metals

• Transition metals are also referred to as d-block elements for this reason.
• In comparison with main group metals, transition metals generally show:

  ⚛ higher densities

  ⚛ higher melting and boiling points

  ⚛ higher ionisation energies

  ⚛ a range of oxidation states

  ⚛ colours in their compounds

  ⚛ the ability to form a wide range of coordination compounds

  ⚛ paramagnetism (ability to attract a magnetic field)

  ⚛ less chemical reactivity than Group 1 (Alkali) metals and Group 2 (Alkali Earth) metals

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Hardness, Density, Melting and Boiling Point of Transition Metals

The table below compares the density, melting point and boiling point of some transition metals with some non-transition metals:

Note that:

• Density of transition metals is greater than the density of the Group 1 and 2 metals.
• Melting point of transition metals is greater than the melting point of Group 1 and 2 metals.
• Boiling point of transition metals is greater than the boiling point of Group 1 and 2 metals.

This is because transition metals have smaller atomic volumes than Group 1 and 2 metals because additional electrons are being progressively added to the inner atomic orbitals resulting in stronger attraction to the nucleus.

These atoms of smaller volume can pack together more closely resulting in higher densities and hardness.

Closer packing results in stronger bonding so more energy is required to melt or boil transition metals.

Ionisation Energy and Chemical Reactivity of the Transition Metals

Ionisation energy (ionization energy) is the energy required to remove an electron form a gaseous atom.

M_(g) → M⁺_(g) + e^-

The table below lists the first ionisation energy of a number of transition metals and of non-transition (main group) metals.

In general, the first ionisation energy of transition metals is higher than for Group 1 and 2 metals.
It requires more energy to remove an electron from a transition metal that it does to remove an electron from a Group 1 or 2 metal.

This is because the smaller atomic radii of transition metals means the valence shell (outer-shell) electrons are more strongly attracted to the nucleus and therefore less easily removed resulting in higher first ionisation energies compared to Group 1 and 2 metals.

Because electrons are less easily lost, the transition metals are less chemically active than Group 1 and 2 metals.

The lower chemical reactivity of the transition metals means they will be placed lower down in the activity series of metals compared to Group 1 and 2 metals.

Since oxidation relates to the loss of electrons, transition metals are less easily oxidised than Group 1 and 2 metals.
Transition metals are therefore weaker reductants than Group 1 and 2 metals.
This is reflected in their standard electrode potentials (E^o) values as shown below:

Oxidation States of Transition Metals

The table below gives the common oxidation states for a number of transition metals and compares them with the oxidation states of non-transition metals:

Note that transition metals generally have more than one oxidation state while Group 1 and 2 metals have only one oxidation state.

This is because the energies of the 3d and 4s orbitals are very close.

Often the lowest oxidation is +2 corresponding to the loss of 2 ns orbital electrons,
where n represents the principal quantum number for the highest energy level.

Higher oxidation states correspond to the additional loss of (n-1)d orbital electrons.

The decrease in maximum states after manganese in the first transition metal series (and after iridium in the second series and osmium in the third series) reflects the difficulty of breaking into a half-filled d subshell.

Coloured Transition Metal Compounds

The table below gives the colour of a number of solid transition metal chlorides as well as the colour of the cation in aqueous solution. This is compared to the colour of non-transition metal chlorides and aqueous solutions containing the cation.

Note the following:

• Solid chlorides of transition metals are often coloured (not white), whereas the solid chlorides of Group 1 and 2 metals are white.
• Aqueous solutions of transition metal chlorides are often coloured (not colourless), whereas the aqueous solutions of Group 1 and 2 chlorides are colourless.
• The colour of the aqueous solution of a transition metal chloride can be different to the colour of the solid transition metal chloride.

A substance will appear coloured if it absorbs light from some portion of the visible spectrum.

Ions with d orbital electrons appear coloured because energy from visible light is absorbed and used to promote a d orbital electron to a higher energy d sublevel (referred to as d-d transitions).

Ions with no d orbital electrons are colourless:

For example, Sc³⁺ and Ti²⁺ are colourless because there are no electrons in d orbitals.

Ions with d¹⁰ electron configurations are colourless, d-d transitions are impossible because the d orbitals are all filled:

For example, Zn²⁺ is colourless because it has 10 electrons occupying the d subshell (all 5 of the d orbitals are full).

Paramagnetism of Transition Metals

Paramagnetism is a weak attraction into a magnetic field.

The table below shows the difference in paramagnetism of transition metals and main group metals:

Note that aqueous solutions of Group 1 and 2 metal cations are not paramagnetic but that aqueous solutions of transition metal cations can be paramagnetic.

Paramagnetism is a weak attraction into a magnetic field.

Substances with unpaired electrons can be paramagnetic.

Paramagnetism is caused by both the orbital and spin motions of electrons (any rotating or revolving charged object generates a magnetic field).

The magnetic fields of paired electrons cancel out, so only unpaired electrons contribute to paramagnetism.

Ferromagnetism and Transition Metals

Ferromagnetism is a strong attraction into a magnetic field.

Ferromagnetism occurs when atoms with unpaired electron spins are just the right distance apart to permit the individual spins to align with each other within a relatively large region.
The individual spins within this region act cooperatively resulting in a large magnetic effect.

Only solids can display the property of ferromagnetism.

The only ferromagnetic elements at room temperature are the following three transition metal elements:

• iron
• cobalt
• nickel

Ferromagnetic compounds such as CrO₂ and Fe₃O₄ also exist.