The first-row transition elements are a series of metals that share several key characteristics. One of these is that they can have more than one oxidation state in their compounds. This is related to the electron configurations of these elements.
Since transition elements can have several different oxidation states, it is possible for their compounds, and the elements themselves, to take part in redox reactions. That is, the lower oxidation states of a transition element can be converted into one of the higher states by oxidation, and its higher oxidation states can be converted into one of the lower states by reduction. These changes necessarily involve the loss and gain of electrons, so they must be redox reactions.
I have shown here how Eꝋred and Eꝋoxid can be used to explain redox reactions. So, it follows that the chemistry of transition elements and their compounds involves electrode potentials.
The first-row transition elements start with titanium and finish with copper. The important redox reactions of six of these elements are the subject of this section, particular mention being paid to the electrode potentials involved.
Vanadium exhibits four common oxidation states, +2, +3, +4 and +5. Under ordinary conditions, vanadium (IV) is considered to be the most stable oxidation state. Vanadium in its lower oxidation states is a reducing agent and in the highest oxidation state it is an oxidizing agent. Therefore the chemistry of vanadium compounds typically involves a conversion from one oxidation state into another.
In the +5 oxidation state, the bonding to the vanadium atom is covalent in character, since it is impossible for a vanadium atom to lose five electrons to form V5+. Probably the most important compound in this oxidation state is vanadium (V) oxide, which is used as a catalyst in the manufacture of sulfuric acid in the contact process.
The aqueous chemistry of vanadium (V) is centered around the vanadyl (V) ion, VO2+ (aq), and the vanadate (V) ion, VO3- (aq). Since vanadium (V) is the highest oxidation state of vanadium, one would expect it to be highly oxidizing, and to be reduced to one of the lower oxidation states. For example, zinc can reduce VO2+ (aq) to V2+(aq) in a series of steps that involves the formation of different coloured solutions as the oxidation state of vanadium changes. Each of these steps can be explained by use of the appropriate electrode potential. The first step is the reduction of vanadyl (V) ion, VO2+(aq). The overall equation can be obtained by combining the two half-equations. The second step is the reduction of vanadyl (IV) ion, VO2+ (aq).
Chromium exhibits three common oxidation states, +2, +3 and +6, of which chromium (III) is the most stable. Chromium is manufactured by the reduction of chromium (III) oxide by a reactive metal (which must be above chromium in the reactivity series). Chromium reacts with dilute acids, such as hydrochloric or sulfuric acid, to give a sky-blue solution that contains the aqueous chromium (II) ion, as indicated by the electrode potential data. The reaction must be carried out in an inert atmosphere to avoid atmospheric oxidation, which would give chromium (III).
Aqueous chromium (III) ions can be oxidized to form chromate (VI). First, they are reacted with aqueous sodium hydroxide to form chromium (III) hydroxide, and then the mixture obtained is boiled with hydrogen peroxide.
The highest oxidation number of a transition element often involves covalent bonding to the transition metal atom. Such is the case with compounds of chromium (VI). These compounds are also very powerful oxidizing agents. Eꝋred for dichromate (VI) has a fairly large positive value, which indicates the ability of dichromate (VI) to oxidize other substances. The equation shows that the product of this reduction is normally the aqueous chromium (II) ion. The use of aqueous dichromate (VI) ion as an oxidizing agent is always associated with the presence of aqueous hydrogen ions (dilute sulfuric acid).
In an alkaline solution, the oxidizing ability is greatly reduced, since dichromate (VI) is converted into chromate (V). Eꝋred for the reduction of the Cr2O72- (aq) ion has a low negative value. The conversion of dichromate (VI) into chromate (VI) is pH dependent. In acid, the equilibrium shifts to the left to form orange dichromate (VI). In akali, the equilibrium shifts to the right to form yellow chromate (V) ion.
Iron (II) ions are oxidized by acidified Cr2O72-(aq) to form Fe3+(aq), and Cr2O72- (aq) is reduced to Cr3+ (aq). The colour changes in this reaction are complicated, since green Fe2+ (aq) reacts with orange Cr2O72- (aq) to form orange Fe3+ (aq) and blue-green Cr3+ (aq). Acidified potassium dichromate (V) can be reduced by zinc in an inert atmosphere to give chromium (II). The reaction must be carried out in an inert atmosphere to prevent oxidation of Cr2+ (aq) to Cr3+ (aq).
Manganese exhibits several common oxidation states, including +2, +4, +6 and +7. The +2 oxidation state is generally considered to be the most stable under normal conditions. Manganese (IV) oxide has an intermediate oxidation number. Therefore, it can act both as a reducing agent and as an oxidizing agent. Manganese (IV) oxide can oxidize hydrochloric acid to form chlorine. It can also be oxidized to give manganate (VI). This is usually achieved by heating the oxide with an oxidizing agent, such as potassium nitrate or potassium chlorate (VI), with potassium hydroxide.
Both the +6 and the +7 oxidation states are highly oxidizing, since the manganese can be reduced to manganese (IV) or manganese (II). Aqueous manganate (VI) is a green solution, which is easily oxidized to give the familiar purple colour of the manganate (VII) ion. The manganate (VII) ion can be prepared directly by reacting manganese (II) ions with oxidizing agents such as sodium bismuthate, NaBiO3, or lead (IV) Oxide. This reaction is used in the estimation of the percentage of manganese in a sample of steel. The steel is reacted with dilute acid to form a solution that contains aqueous iron (II) ions and aqueous manganese (II) ions are oxidized to give purple manganate (VII). The concentration of the manganate (VI) can be determined using ultraviolet-visible spectroscopy.
Aqueous potassium manganate (VII) is used in laboratories as an oxidizing agent for organic preparations, volumetric analysis and qualitative analysis. Acidified potassium manganite (VII) is used to test for reducing agents: it reacts to form manganese (II) ions and the distinct colour of the manganite (VII) ions disappears. For example, iodide ions reduce acidified manganite (VII) ions to form manganese (II) ions and iodine.
In volumetric analysis, acidified potassium manganate (VII) is used to determine the concentration of reducing agents. Since it changes colour during the titration, there is no need to have an indicator. If potassium manganate (VII) is used in the burette, the end-point is the first appearance of a purple-pink colour. A typical example would be the reaction of aqueous iron (II) ions with acidified potassium manganate (VII).
The lithium-manganese (IV) oxide cell – popularly referred to as the lithium cell – cannot have a water-based electrolyte because the lithium would react with the water to form hydrogen. The great advantage of using lithium as one of the electrodes is that its oxidation potential is very large, which gives the battery a high voltage. The battery can therefore deliver a low current for a long time, which makes it an ideal power supply for such items as watches and heart pacemakers.
Iron exhibits two common oxidation states, +2 and +3, both of which are ionic. Iron does form another oxidation state, +6, but it occurs in only a few compounds. As the electrode potentials suggest, iron reacts with dilute acids, such as sulfuric acid, to form iron (II) salts. The iron (II) ions are very susceptible to aerial oxidation to form iron (III) ions. The stability of these two oxidation states is very much dependent on the pH. In acidic conditions, aqueous iron (II) ions are oxidized to give iron (III) ions. In alkaline conditions, the oxidation is even more favourable:
Any unpainted iron object shows evidence of rusting. The surface of iron forms a flaky orange solid, the composition of which is best described as hydrated iron (III) oxide. Rusting is, in fact, a complicated electrochemical process that occurs on the surface of iron. When iron is in contact with a drop of water, a redox reaction occurs:
Oxidation half-equation: Fe(s) → Fe2+(aq) + 2e-
Reduction half-equation: O2 (g) + 2H2O(l) + 4e- → 4OH- (aq)
These two reactions take place in different areas of the iron, which results in the formation of an anodic region and a cathodic region. The electrons move through the metal from the anodic region to the cathodic region. The circuit is completed by ions that move through the water. Without the water, the circuit is not complete and rusting cannot take place. Essentially, the water is acting as a salt bridge. If the water contains electrolytes, such as sodium chloride, then the concentration of ions in the droplet is higher and so the rate of rusting increases. The iron (II) ions and the hydroxide in the water droplet are precipitated as iron (II) hydroxide, which is further oxidized to form hydrated iron (III) oxide or rust.
Any layer that is impervious to water protects iron from rusting. However, the protection provided by paints and other widely used materials is of limited duration, because of such factors as physical damage, chemical deterioration and weathering. So, alternative methods of rust protection use electrochemical principles instead. One of the best known involves giving the iron a protective layer of zinc by a process called galvanization, which is described below.
Iron pipes and tanks in the ground are often protected by a block of magnesium or zinc. The block of magnesium (or zinc) is attached to the iron object and, since magnesium (or zinc) has a higher positive oxidation potential than iron, it oxidizes in preference (sacrificially) to the iron. The iron object acts as the cathode in this method.
Tin-plated cans used in the food industry provide another example of electrochemical protection. The oxidation potential for tin has a lower positive value than that for iron. Therefore, tin is less likely to react with moist oxygen and so protects the surface of the iron. There is one drawback with tin plating. It is easily scratched to reveal the iron, in which case the can rusts very rapidly. What happens is that once the iron is in contact with moist air, it gives sacrificial protection to the less reactive tin.
As already stated, a coating of zinc protects iron from rusting. Zinc-coated iron is known as galvanized iron. Galvanized iron is protected mainly because, in the electrochemical cell formed by galvanization, the zinc is preferentially oxidized. The oxidation potential for zinc has a larger positive value than that for iron.
However, were it not for another reaction taking place at the same time, all the zinc would eventually be oxidized and rusting could start. The zinc hydroxide produced reacts with carbon dioxide in the air to form a layer of a zinc hydroxide-zinc carbonate compound that adheres firmly to the iron to give further protection.
