Variability and relative stability of oxidation states
Variability and relative stability of oxidation states
In this essay the concentration will be on the oxidation states of the transition metals rather than the entire periodic table.
An oxidation state is a number assigned to an element in a compound according to some rules. This number enables us to describe oxidation-reduction reactions, and balancing redox chemical reactions. Compared to Group II elements such as calcium, transition elements form ions with a wide variety of oxidation states. Calcium ions typically don't lose more than two electrons, whereas transition metals can lose up to nine.
The energies required to remove electrons from calcium are low until you try to remove electrons from below its outer two s orbitals. In fact Ca3+ has an ionisation enthalpy so high that it rarely occurs naturally. However a transition element like vanadium has roughly linear increasing ionisation enthalpies throughout its s and d orbitals, due to the close energy difference between the 3d and 4s orbitals. Transition metal ions are therefore commonly found in very high states.
Certain patterns can be seen to emerge across the period of transition elements:
The number of oxidation states of each ion increases up to Mn, after which they start to drop. This drop is due to the stronger pull from the protons in the nucleus towards the electrons, making them harder to remove.
When the elements are in lower oxidation states, they can be found as simple ions. However elements in higher oxidation states are usually bonded covalently to electronegative compounds such as O or F, often in an anion.
Properties with respect to the stability of oxidation states:
Higher oxidation state ions become less stable across the period.
Ions in higher oxidation states tend to make good oxidising agents, whereas elements in low oxidation states become reducing agents.
The 2+ ions across the period start as strong reducing agents, and become more stable.
This is a preview of the whole essay
The 3+ ions start stable and become more oxidising across the period.
One point about the oxidation states of transition metals deserves particular attention: Transition-metal ions with charges larger than +3 cannot exist in aqueous solution. Consider the following reaction in which manganese is oxidized from the +2 to the +7 oxidation state.
When the manganese atom is oxidized, it becomes more electronegative. In the +7 oxidation state, this atom is electronegative enough to react with water to form a covalent oxide, MnO4-. Mn(VII) is not the only example of an oxidation state powerful enough to decompose water. As soon as Mn2+ is oxidized to Mn(IV), it reacts with water to form MnO2. Vanadium exists in aqueous solutions as the V2+ ion. But once it is oxidized to the +4 or +5 oxidation state, it reacts with water to form the VO2+ or VO2+ ion. The Cr3+ ion can be found in aqueous solution. But once this ion is oxidized to Cr(VI), it reacts with water to form the CrO42- and Cr2O72- ions.
The table below summarises known oxidation numbers of the first row transition elements.
Known Oxidation Numbers of First row Transition Elements*
* The oxidation number zero usually assigned to elemental state has been omitted from the table. The elements Cr to Co form several metal carbonyl compounds where the metals are in zero oxidation state.
From this table it is found that:
There is an increase in the number of oxidation states from Sc to Mn. All seven oxidation states are exhibited by Mn. The formal oxidation number of +7 represents the formal loss of all seven electrons from 3d and 4s orbitals. In fact all of the elements in the series can utilize all the electrons in their 3d and 4s orbitals. There is a decrease in the number of oxidation states from Mn to Zn. This is because the pairing of d-electrons occurs after Mn (Hund's rule) which in turn decreases the number of available unpaired electrons and hence, the number of oxidation states. The stability of higher oxidation states decreases in moving from Sc to Zn. Mn(VII) and Fe(VI) are powerful oxidizing agents and the higher oxidation states of Co, Ni and Zn are unknown. The relative stability of +2 state with respect to higher oxidation states, particularly +3 state increases in moving from left to right. This is justifiable since it will be increasingly difficult to remove the third electron from the d orbital. There is a tendency of intermediate oxidation states to disproportionate. As for example,
Mn(VI)→ Mn(IV) + Mn(VII)
Cu(I) → Cu(0) + Cu(II).
The lower oxidation states are usually found in ionic compounds and higher oxidation states tend to be involved in covalent compounds.
The relative stability of oxidation states is an extremely important topic in transition metal chemistry and is usually discussed in terms of standard reduction potential (E°) values. Thermodynamically E° values are equated to ΔG° values in the form of the well known relationship: ΔG° = -nFE° where n = number of electrons involved and F = Faraday constant. Hence, the E° values indicate the possibility of spontaneous change from one oxidation state to the other. Predictions about the stability of an oxidation state of an element can be made from the tables of Redox values shown below.
A Frost diagram for the first series of d-block elements in acidic solution (pH = 0).
The bold numbers designate the group numbers and the broken line connects species in their group oxidation states.
The group oxidation number can be achieved in elements that lie toward the left of the d-block but not on the right i.e. Group 3 elements are found in aqueous solution only with an oxidation number of +3. The oxidation number is never achieved after group 9. This maximum oxidation number corresponds with the increase of noble character from left to right across the d-block. The Frost diagram above, shows that scandium, titanium and vanadium fall in the lower part of the diagram, which means that the metal and any intermediate species are readily oxidised to the group oxidation state. Species in the upper part of the diagram, such as manganese (+7) and chromium (+6), are readily reduced. Chromate (CrO42-), permanganate (MnO4-) and ferrate (FeO42-) are strong oxidising agents which increase in strength from chromate to ferrate. This trend is another illustration of the decreasing stability of the maximum attainable oxidation state for groups 6, 7 and 8.
In groups 4 to 10 the highest oxidation state of an element becomes more stable on descending a group, with the greatest change in stability occurring between the first two rows of the d-block. This trend is illustrated in the Frost diagram on the next page. The increasing stability of high oxidation states for the heavier d-block metals in the formulas of their halides is shown in the table below:
Highest oxidation state d-block binary halides*
*The formulas show the least electronegative halide that brings out the highest oxidation state
†CrF6 exists for several days at room temperature in a passivated Monel container
The hexafluorides of the heavier d-block elements have been prepared from group 6 to 10 (as in PtF6). In keeping with the stability of high oxidation states for the heavier metals, WF6 is not a significant oxidising agent. However, the oxidising character of the hexafluorides increases to the right, and PtF6 is so strong that it can oxidise O2 to O2+.
A Frost diagram for the chromium group in the d-block in acidic solution (pH = 0)
Most +1 d-block metal cations disproportionate (to M and M2+) because the bonds in the solid metal are so strong. However the more noble metals, copper, silver and gold, form many salts containing M+. For other d metals the +2 oxidation state is generally the lowest one to consider in aqueous solution and in combination with hard ligands. Many of the dipositive aqua ions, M2+, are coloured as a result of d-d transitions in the visible region of the spectrum. For example, Mn2+(aq) is pale pink, Fe2+ (aq) is pale green, Co2+(aq) is pink, Ni2+(aq) is green and Cu2+(aq) is blue. The +2 oxidation state becomes increasingly common from left to right across the period. For example, Sc2+(aq) (Group 3) is unknown and Ti2+(aq) (Group 4) is not readily accessible. For groups 5 and 6, V2+(aq) and Cr2+(aq) are thermodynamically unstable with respect to oxidation by H+. Past chromium (for Mn2+, Fe2+, Co2+, Ni2+, and Cu2+), the +2 state is stable with respect to reaction with water, and only Fe2+ is oxidised by air. The trend toward increased stability of the lower oxidation states in the middle and the right of the d-block can be understood from the general increase in ionisation energies from left to right across a period in the d-block.
In contrast to the 3d metals, the 4d and 5d only rarely form simple M2+(aq) ions. M(II) complexes with σ-donor ligands are common for the 3d metals but M(II) complexes of 4d and 5d metals are less common. They generally contain π-acceptor ligands.
In conclusion it is found that the transition elements show a wide range of oxidation states which show a differing range of stabilities.
D.F. Shriver, P.W. Atkins; Inorganic chemistry; Third edition
G. Wulfsberg; Inorganic chemistry
P.W. Atkins, L. Jones; Chemical principles; Second edition.
Here's what a teacher thought of this essay
This is a very detailed account of transition metals, including advanced work for even A2 level. It is set out in a logical and coherent manner and would be very interesting to someone researching transition metals. This piece of work is 5 stars out of 5.