Transition Metals: Variable Oxidation States

Topic 15: Transition Metals 

1. In order to develop their practical skills, students will be encouraged to carry out a range of practical experiments related to this topic. Possible experiments to be discussed include the stepwise reduction of vanadium(V) to vanadium(II), investigating the reactions of copper(II) ions or chromium(III) ions, using sodium hydroxide and ammonia solution to identify transition metal ions, investigating autocatalysis and preparing a complex transition metal salt.
2. Mathematical skills that will be developed in this topic include investigating the geometry of different transition metal complexes.
3. Within this topic, students will consider the model for the filling of electron orbitals encountered earlier in their course, and see how limitations in that model indicate the need for more sophisticated explanations. They will also appreciate that catalyst research is a frontier area, and one that provides an opportunity to show how the scientific community reports and validates new knowledge.

Edexcel A level Chemistry (2017)

Topic 15A: Principles of transition metal chemistry

15A/3 To be able to understand why transition metals show variable oxidation number

What is oxidation state?

Oxidation state is the oxidation number a particular species carries.

So the sodium ion Na⁺ has oxidation number +1 and sodium is in oxidation state +1.

Or the chloride ion Cl^— has oxidation number —1 and chlorine is in oxidation state —1.

Transition metals show variable oxidation states unlike group 1 or Group 17 elements.

The illustration below shows these oxidation states some of which of more stable (shown in colour) at room temperature than others. 

Variable Oxidation States 

Question is why do these metals exhibit variable oxidation states?

The answer has to do with the fact that the 4s energy level and the 3d energy level are very close in value so that the metals can bond to other elements using electrons from both the 4s and the 3d levels. 

Here are the 3d and 4s electron arrangements for the stable oxidation states of these metals

So take the very stable purple manganate(VII) ion (MnO₄^—) the manganese ion (oxidation state 7+) engages in 4 covalent bonds with the oxide ions (oxidation state —2) giving rise to the tetrahedral structured ion overall charge —1 (see below).

This ion also illustrates another feature of the higher oxidation states of the transition metals which is that they are stabilized as oxo anions i.e the metal ion is combined with oxide ions.

Here are a few common examples you will have come across in your college or A level chemistry studies:

Orange Dichromate (VI)           Cr₂O₇^2—

Yellow Chromate (VI)               CrO₄^2—

Black Manganese (IV)oxide       MnO₂

Yellow Metavanadate (V)          VO₃^—

Purple Manganate(VII)             MnO₄^—

Green Manganate(VI)               MnO₄^2—

Transition Metals: Electron Arrangements

Topic 15: Transition Metals

In order to develop their practical skills, students will be encouraged to carry out a range of practical experiments related to this topic. Possible experiments to be discussed include the stepwise reduction of vanadium(V) to vanadium(II), investigating the reactions of copper(II) ions or chromium(III) ions, using sodium hydroxide and ammonia solution to identify transition metal ions, investigating autocatalysis and preparing a complex transition metal salt.

Mathematical skills that will be developed in this topic include investigating the geometry of different transition metal complexes.

Within this topic, students will consider the model for the filling of electron orbitals encountered earlier in their course, and see how limitations in that model indicate the need for more sophisticated explanations. They will also appreciate that catalyst research is a frontier area, and one that provides an opportunity to show how the scientific community reports and validates new knowledge.

Edexcel A level Chemistry (2017)

Topic 15A: Principles of transition metal chemistry

Here are the learning objectives relating to the d block:

15A/1: To be able to deduce the electronic configurations of atoms and ions of the d-block elements of Period 4 (Sc–Zn), given the atomic number and charge (if any).

15A/2: To know that transition metals are d-block elements that form one or more stable ions with incompletely-filled d-orbitals.

The Periodic Table: the d block

Here is the d block:

In this topic, we begin by discussing how to work out the electron configurations of the first row of d block elements then the vexed question (certainly in the minds of some teachers of chemistry) as to what really constitutes a transition metal and whether all first row d block elements are also transition metals.

But first here are some pictures of these metals:

Determining the electron configurations of the first row d block elements.

Starting with the atomic numbers of these d block elements, you ought to be able to correctly build up the electron configurations of these elements.

Up to Argon Z=18, all the metals have the same electron configuration.  Things only get tricky beyond that number.  For example, iron Z= 26 so as the 4s fills before the 3d that leaves 6 electrons in the 3d and an electron configuration of :

Fe: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²

But look at Chromium Z=24 how do the six electrons fill the 3d and 4s?

Here we remember that half filled subshells confer an added degree of energetic stability to the atom so chromium has 3d⁵ 4s¹

Cr: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹

By a similar consideration the electron configuration of copper Cu Z= 29 the extra 11 electrons fill the 3d and 4s, one in the 4s and the rest (10) in the 3d to confer the greatest energetic stability. 

Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹

So remember that the 4s fills before the 3d and that the electrons above the 18 fill the subshells to confer the greatest energetic stability by half filling or filling subshells. 

Definition of the term Transition metal

You will see from the two learning objectives above from the British Edexcel A level course that a d block element might not also be a transition metal.

Let’s look first at the electron configurations of all the first row d-block elements.

Here they are easy to find on the Internet. 

You’ll find them in at least two representations: either as electrons in boxes (as arrows) or with subshell electrons as superscripts.

Electrons in boxes has the advantage of showing the opposite electron spins and keeping the shells together i.e. all third shell subshells are together.

Now which of these d block elements are said to be transition metals?

Look at the electron configurations and if the atom of the element or its ion has an incomplete d subshell then we are going to call those elements transition elements.

The result is that zinc is not a transition metal but it is a d block element. 

Zinc atom:  1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s²

And the zinc ion Zn²⁺:  1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰

In both examples the 3d subshell is full with 10 electrons so the metal is not a transition element according to the Edexcel definition.

So what is your verdict on Scandium: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹ 4s²

when its Sc³⁺ ion is: 1s² 2s² 2p⁶ 3s² 3p⁶   

since this ion has no 3d orbital at all?

According to the definition given above it too is not a transition metal since its ion has no 3d subshell.

Effect of half filled orbitals:

Another thing you should see from the table of electron configurations is the transition metals with half filled d subshells.

Note first the electron configuration of chromium:

Cr: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹

According to Hund’s Rule, the subshells half fill then the electrons in the orbitals start to pair up.  So chromium is 3d⁵ 4s¹ and both are half full.

Second, note the electron configurations of Mn²⁺ and Fe³⁺ these are stable oxidation states of these metals because the electron configuration in each case is 3d⁵ that is an energetically stable half filled 3d subshell.

Mn:  1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s²

Mn²⁺:  1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵

And

Fe:  1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²

Fe³⁺:  1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵

Third, note the electron configuration of copper.  It has a half filled 4s orbital and a full 3d subshell.

So what we find is that Cu⁺ ions are stable and colourless because the 3d subshell is full but as soon as the Cu⁺ is oxidised to Cu²⁺ the 3d subshell becomes 3d⁹ and the ions are energetically stable and blue. 

Cu: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s¹

Cu⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰  colourless ions

Cu²⁺: 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁹  blue ions

The next post will begin to explain the individual properties of d block elements.