Chemistry » Transition Metals » Occurrence, Preparation, and Properties of Transition Metals and Their Compounds

# Transition Metal Compounds

## Transition Metal Compounds

The bonding in the simple compounds of the transition elements ranges from ionic to covalent. In their lower oxidation states, the transition elements form ionic compounds; in their higher oxidation states, they form covalent compounds or polyatomic ions. The variation in oxidation states exhibited by the transition elements gives these compounds a metal-based, oxidation-reduction chemistry. The chemistry of several classes of compounds containing elements of the transition series follows.

## Halides

Anhydrous halides of each of the transition elements can be prepared by the direct reaction of the metal with halogens. For example:

$$\text{2Fe}(s)+{\text{3Cl}}_{2}(g)\;⟶\;{\text{2FeCl}}_{3}(s)$$

Heating a metal halide with additional metal can be used to form a halide of the metal with a lower oxidation state:

$$\text{Fe}(s)+{\text{2FeCl}}_{3}(s)\;⟶\;{\text{3FeCl}}_{2}(s)$$

The stoichiometry of the metal halide that results from the reaction of the metal with a halogen is determined by the relative amounts of metal and halogen and by the strength of the halogen as an oxidizing agent. Generally, fluorine forms fluoride-containing metals in their highest oxidation states. The other halogens may not form analogous compounds.

In general, the preparation of stable water solutions of the halides of the metals of the first transition series is by the addition of a hydrohalic acid to carbonates, hydroxides, oxides, or other compounds that contain basic anions. Sample reactions are:

$${\text{NiCO}}_{3}(s)+\text{2HF}(aq)\;⟶\;{\text{NiF}}_{2}(aq)+{\text{H}}_{2}\text{O}(l)+{\text{CO}}_{2}(g)$$

$$\text{Co}{(\text{OH})}_{2}(s)+\text{2HBr}(aq)\;⟶\;{\text{CoBr}}_{2}(aq)+\;{\text{2H}}_{2}\text{O}(l)$$

Most of the first transition series metals also dissolve in acids, forming a solution of the salt and hydrogen gas. For example:

$$\text{Cr}(s)+\text{2HCl}(aq)\;⟶\;{\text{CrCl}}_{2}(aq)+{\text{H}}_{2}(g)$$

The polarity of bonds with transition metals varies based not only upon the electronegativities of the atoms involved but also upon the oxidation state of the transition metal. Remember that bond polarity is a continuous spectrum with electrons being shared evenly (covalent bonds) at one extreme and electrons being transferred completely (ionic bonds) at the other.

No bond is ever 100% ionic, and the degree to which the electrons are evenly distributed determines many properties of the compound. Transition metal halides with low oxidation numbers form more ionic bonds. For example, titanium(II) chloride and titanium(III) chloride (TiCl2 and TiCl3) have high melting points that are characteristic of ionic compounds, but titanium(IV) chloride (TiCl4) is a volatile liquid, consistent with having covalent titanium-chlorine bonds. All halides of the heavier d-block elements have significant covalent characteristics.

The covalent behavior of the transition metals with higher oxidation states is exemplified by the reaction of the metal tetrahalides with water. Like covalent silicon tetrachloride, both the titanium and vanadium tetrahalides react with water to give solutions containing the corresponding hydrohalic acids and the metal oxides:

$${\text{SiCl}}_{4}(l)+2{\text{H}}_{2}\text{O}(l)\;⟶\;{\text{SiO}}_{2}(s)+\text{4HCl}(aq)$$

$${\text{TiCl}}_{4}(l)+2{\text{H}}_{2}\text{O}(l)\;⟶\;{\text{TiO}}_{2}(s)+\text{4HCl}(aq)$$

## Oxides

As with the halides, the nature of bonding in oxides of the transition elements is determined by the oxidation state of the metal. Oxides with low oxidation states tend to be more ionic, whereas those with higher oxidation states are more covalent. These variations in bonding are because the electronegativities of the elements are not fixed values.

The electronegativity of an element increases with increasing oxidation state. Transition metals in low oxidation states have lower electronegativity values than oxygen; therefore, these metal oxides are ionic. Transition metals in very high oxidation states have electronegativity values close to that of oxygen, which leads to these oxides being covalent.

The oxides of the first transition series can be prepared by heating the metals in air. These oxides are Sc2O3, TiO2, V2O5, Cr2O3, Mn3O4, Fe3O4, Co3O4, NiO, and CuO.

Alternatively, these oxides and other oxides (with the metals in different oxidation states) can be produced by heating the corresponding hydroxides, carbonates, or oxalates in an inert atmosphere. Iron(II) oxide can be prepared by heating iron(II) oxalate, and cobalt(II) oxide is produced by heating cobalt(II) hydroxide:

$${\text{FeC}}_{2}{\text{O}}_{4}(s)\;⟶\;\text{FeO}(s)+\text{CO}(g)+{\text{CO}}_{2}(g)$$

$${\text{Co(OH)}}_{2}(s)\;⟶\;\text{CoO}(s)+{\text{H}}_{2}\text{O}(g)$$

With the exception of CrO3 and Mn2O7, transition metal oxides are not soluble in water. They can react with acids and, in a few cases, with bases. Overall, oxides of transition metals with the lowest oxidation states are basic (and react with acids), the intermediate ones are amphoteric, and the highest oxidation states are primarily acidic. Basic metal oxides at a low oxidation state react with aqueous acids to form solutions of salts and water. Examples include the reaction of cobalt(II) oxide accepting protons from nitric acid, and scandium(III) oxide accepting protons from hydrochloric acid:

$$\text{CoO}(s)+{\text{2HNO}}_{3}(aq)\;⟶\;\text{Co}{({\text{NO}}_{3})}_{2}(aq)+{\text{H}}_{2}\text{O}(l)$$

$${\text{Sc}}_{2}{\text{O}}_{3}(s)+\text{6HCl}(aq)\;⟶\;{\text{2ScCl}}_{3}(aq)+{\text{3H}}_{2}\text{O}(l)$$

The oxides of metals with oxidation states of 4+ are amphoteric, and most are not soluble in either acids or bases. Vanadium(V) oxide, chromium(VI) oxide, and manganese(VII) oxide are acidic. They react with solutions of hydroxides to form salts of the oxyanions $${\text{VO}}_{4}{}^{3-},$$$${\text{CrO}}_{4}{}^{2-},$$ and $${\text{MnO}}_{4}{}^{-}.$$ For example, the complete ionic equation for the reaction of chromium(VI) oxide with a strong base is given by:

$${\text{CrO}}_{3}(s)+{\text{2Na}}^{\text{+}}(aq)+{\text{2OH}}^{\text{−}}(aq)\;⟶\;{\text{2Na}}^{\text{+}}(aq)+{\text{CrO}}_{4}{}^{2-}(aq)+{\text{H}}_{2}\text{O}(l)$$

Chromium(VI) oxide and manganese(VII) oxide react with water to form the acids H2CrO4 and HMnO4, respectively.

## Hydroxides

When a soluble hydroxide is added to an aqueous solution of a salt of a transition metal of the first transition series, a gelatinous precipitate forms. For example, adding a solution of sodium hydroxide to a solution of cobalt sulfate produces a gelatinous pink or blue precipitate of cobalt(II) hydroxide. The net ionic equation is:

$${\text{Co}}^{2+}(aq)+{\text{2OH}}^{\text{−}}(aq)\;⟶\;\text{Co}{(\text{OH})}_{2}(s)$$

In this and many other cases, these precipitates are hydroxides containing the transition metal ion, hydroxide ions, and water coordinated to the transition metal. In other cases, the precipitates are hydrated oxides composed of the metal ion, oxide ions, and water of hydration:

$${\text{4Fe}}^{3+}(aq)+{\text{6OH}}^{\text{−}}(aq)+\text{n}\;{\text{H}}_{2}\text{O}(l)\;⟶\;2{\text{Fe}}_{2}{\text{O}}_{3}\text{·}(\text{n}+3){\text{H}}_{2}\text{O}(s)$$

These substances do not contain hydroxide ions. However, both the hydroxides and the hydrated oxides react with acids to form salts and water. When precipitating a metal from solution, it is necessary to avoid an excess of hydroxide ion, as this may lead to complex ion formation as discussed later in this tutorial. The precipitated metal hydroxides can be separated for further processing or for waste disposal.

## Carbonates

Many of the elements of the first transition series form insoluble carbonates. It is possible to prepare these carbonates by the addition of a soluble carbonate salt to a solution of a transition metal salt. For example, nickel carbonate can be prepared from solutions of nickel nitrate and sodium carbonate according to the following net ionic equation:

$${\text{Ni}}^{2+}(aq)+{\text{CO}}_{3}{}^{2-}\;⟶\;{\text{NiCO}}_{3}(s)$$

The reactions of the transition metal carbonates are similar to those of the active metal carbonates. They react with acids to form metals salts, carbon dioxide, and water. Upon heating, they decompose, forming the transition metal oxides.

## Other Salts

In many respects, the chemical behavior of the elements of the first transition series is very similar to that of the main group metals. In particular, the same types of reactions that are used to prepare salts of the main group metals can be used to prepare simple ionic salts of these elements.

A variety of salts can be prepared from metals that are more active than hydrogen by reaction with the corresponding acids: Scandium metal reacts with hydrobromic acid to form a solution of scandium bromide:

$$\text{2Sc}(s)+\text{6HBr}(aq)\;⟶\;2{\text{ScBr}}_{3}(aq)+{\text{3H}}_{2}(g)$$

The common compounds that we have just discussed can also be used to prepare salts. The reactions involved include the reactions of oxides, hydroxides, or carbonates with acids. For example:

$$\text{Ni}{(\text{OH})}_{2}(s)+2{\text{H}}_{3}{\text{O}}^{\text{+}}(aq)+{\text{2ClO}}_{4}{}^{\text{−}}(aq)\;⟶\;{\text{Ni}}^{2+}(aq)+{\text{2ClO}}_{4}{}^{\text{−}}(aq)+{\text{4H}}_{2}\text{O}(l)$$

Substitution reactions involving soluble salts may be used to prepare insoluble salts. For example:

$${\text{Ba}}^{2+}(aq)+{\text{2Cl}}^{\text{−}}(aq)+{\text{2K}}^{\text{+}}(aq)+{\text{CrO}}_{4}{}^{2-}(aq)\;⟶\;{\text{BaCrO}}_{4}(s)+{\text{2K}}^{\text{+}}(aq)+{\text{2Cl}}^{\text{−}}(aq)$$

In our discussion of oxides in this section, we have seen that reactions of the covalent oxides of the transition elements with hydroxides form salts that contain oxyanions of the transition elements.

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