8 julio, 2024

Hydroxides: what they are, properties, nomenclature and examples

What are hydroxides?

The hydroxides are inorganic and ternary compounds that consist of the interaction between a metal cation and the functional group OH (hydroxide anion, OH–). Most of them are ionic in nature, although they can also have covalent bonds.

For example, a hydroxide can be represented as the electrostatic interaction between the M+ cation and the OH– anion, or as the covalent union through the M-OH bond (bottom image). In the first the ionic bond is given, while in the second the covalent. This fact essentially depends on the metal or cation M+, as well as its charge and ionic radius.

Since most of them come from metals, it is equivalent to mention them as metal hydroxides.

How are hydroxides formed?

There are two main synthetic routes: through the reaction of the corresponding oxide with water, or with a strong base in an acid medium:

MO + H2O => M(OH)2

MO + H+ + OH– => M(OH)2

Only those metal oxides soluble in water react directly to form the hydroxide (first chemical equation). Others are insoluble and require acidic species that release M+, which then interacts with the OH– coming from the strong bases (second chemical equation).

However, such strong bases are metal hydroxides NaOH, KOH and others from the group of alkali metals (LiOH, RbOH, CsOH). These are ionic compounds that are highly soluble in water, therefore their OH– are free to participate in chemical reactions.

On the other hand, there are insoluble metal hydroxides and consequently they are very weak bases. Some of them are even acids, as is the case with telluric acid, Te(OH)6.

The hydroxide establishes a solubility equilibrium with the solvent around it. If it is water, for example, then the equilibrium is expressed as follows:

M(OH)2 <=> M2+(aq) + OH–(aq)

Where (ac) denotes that the medium is aqueous. When the solid is insoluble, the dissolved OH concentration is small or negligible. For this reason, insoluble metal hydroxides cannot generate solutions as basic as those of NaOH.

From the above it can be deduced that the hydroxides exhibit very different properties, linked to the chemical structure and the interactions between the metal and the OH. Thus, although many are ionic, with varied crystalline structures, others instead present complex and disordered polymeric structures.

Properties of hydroxides

OH– anion

The hydroxyl ion is an oxygen atom covalently bonded to a hydrogen. Thus, this can easily be represented as OH–. The negative charge is placed on the oxygen, making this anion an electron donating species: a base.

If OH– donates its electrons to a hydrogen, an H2O molecule is formed. It can also donate its electrons to positively charged species: like the M+ metal centers. Thus, a coordination complex is formed through the dative bond M–OH (oxygen provides the pair of electrons).

However, for this to happen, the oxygen must be able to coordinate efficiently with the metal, otherwise the interactions between M and OH will have a strong ionic character (M+ OH–).

Since the hydroxyl ion is the same in all hydroxides, the difference between all of them then lies in the cation that accompanies it.

Also, because this cation can come from any metal in the periodic table (groups 1, 2, 13, 14, 15, 16, or transition metals), the properties of such hydroxides vary enormously, although they all contemplate in common some aspects.

ionic and basic character

Hydroxides, although they have coordination bonds, have a latent ionic character. In some, such as NaOH, its ions are part of a crystal lattice made up of Na+ cations and OH– anions in 1:1 ratios; that is, for every Na+ ion there is a counterpart OH– ion.

Depending on the charge of the metal, there will be more or less OH– anions around it. For example, for a metal cation M2+ there will be two OH– ions interacting with it: M(OH)2, which is sketched as HO– M2+ OH–.

In the same way it happens with M3+ metals and with others with more positive charges (although they rarely exceed 3+).

This ionic character is responsible for many of the physical properties, such as melting and boiling points. These are high, reflecting the electrostatic forces operating within the crystal lattice. Likewise, when hydroxides dissolve or melt they can conduct electrical currents due to the mobility of their ions.

However, not all hydroxides present the same crystalline networks. Those with the most stable will be less likely to dissolve in polar solvents like water. As a general rule, the more disparate the ionic radii of M+ and OH– are, the more soluble they are.

periodic trend

This explains why the solubility of alkali metal hydroxides increases as you go down the group. Thus, the increasing order of solubilities in water for these is as follows: LiOH

OH– is a small anion, and as the cation becomes bulkier, the lattice becomes energetically weaker.

On the other hand, alkaline earth metals form less soluble hydroxides due to their greater positive charges. This is because M2+ attracts OH– more strongly than M+. Likewise, its cations are smaller, and therefore less unequal in size with respect to OH–.

The result of this is experimental evidence that NaOH is much more basic than Ca(OH)2. The same reasoning can be applied to other hydroxides, either for those of the transition metals, or for those of the p-block metals (Al, Pb, Te, etc.).

Also, the smaller and larger the ionic radius and positive charge of M+, the lower the ionic character of the hydroxide, in other words, those with very high charge densities. An example of this occurs with beryllium hydroxide, Be(OH)2. Be2+ is a very small cation and its divalent charge makes it electrically very dense.

amphotericism

Hydroxides M(OH)2 react with acids to form an aquacomplex, that is, M+ ends up surrounded by water molecules. However, there are a limited number of hydroxides that can also react with bases. These are what are known as amphoteric hydroxides.

Amphoteric hydroxides react with both acids and bases. The second situation can be represented by the following chemical equation:

M(OH)2 + OH– => M(OH)3–

But how to determine if a hydroxide is amphoteric? Through a simple laboratory experiment. Because many metal hydroxides are insoluble in water, adding a strong base to a solution with dissolved M+ ions, eg Al3+, will precipitate the corresponding hydroxide:

Al3+(aq) + 3OH–(aq) => Al(OH)3(s)

But since there is an excess of OH– the hydroxide continues to react:

Al(OH)3(s) + OH– => Al(OH)4–(aq)

As a result, the new negatively charged complex is solvated by the surrounding water molecules, dissolving the white solid aluminum hydroxide. Those hydroxides that remain unchanged with the extra addition of base do not behave as acids and, therefore, are not amphoteric.

structures

Hydroxides can have crystal structures similar to those of many salts or oxides; some simple, and others very complex. In addition, those where there is a decrease in the ionic character may present metallic centers linked by oxygen bridges (HOM–O–MOH).

In solution the structures are different. Although for highly soluble hydroxides it is enough to consider them as ions dissolved in water, for others it is necessary to take coordination chemistry into account.

Thus, each M+ cation can coordinate to a limited number of species. The larger it is, the greater the number of water or OH– molecules bound to it. From here arises the famous coordination octahedron of many metals dissolved in water (or in any other solvent): M(OH2)6+n, where n is equal to the positive charge of the metal.

Cr(OH)3, for example, actually forms an octahedron. As? Considering the compound as [Cr(OH2)3(OH)3]of which three of the water molecules are replaced by OH– anions.

If all the molecules were replaced by OH–, then the complex with a negative charge and an octahedral structure would be obtained. [Cr(OH)6]3-. The charge -3 is the result of the six negative charges of the OH–.

dehydration reaction

The hydroxides can be considered as «hydrated oxides». However, in them «water» is in direct contact with M+; while in the hydrated oxides MO·nH2O, the water molecules are part of an external coordination sphere (they are not close to the metal).

Such water molecules can be extracted by heating a hydroxide sample:

M(OH)2 + Q(heat) => MO + H2O

MO is the metal oxide formed as a result of hydroxide dehydration. An example of this reaction is the one observed when cupric hydroxide, Cu(OH)2, is dehydrated:

Cu(OH)2 (blue) + Q => CuO (black) + H2O

Hydroxide nomenclature

What is the proper way to mention hydroxides? The IUPAC raised three nomenclatures for this purpose: the traditional, the stock and the systematic. It is correct to use any of the three, however, for some hydroxides it may be more comfortable or practical to mention it in one way or another.

Traditional

The traditional nomenclature consists simply of adding the suffix -ico to the highest valence that the metal presents; and the suffix -oso to the lowest. So, for example, if the metal M has valences +3 and +1, the hydroxide M(OH)3 will be called hydroxide (name of the metal)icowhile MOH hydroxide (metal name)bear.

To determine the valency of the metal in the hydroxide, just look at the number after the OH enclosed in parentheses. Thus, M(OH)5 means that the metal has a charge or valence of +5.

The main drawback of this nomenclature, however, is that it can be complicated for metals with more than two oxidation states (such as chromium and manganese). For such cases, the prefixes hyper- and hypo- are used to denote the highest and lowest valences.

Thus, if M, instead of having only +3 and +1 valences, also has +4 and +2, then the names of its higher and lower valence hydroxides are: hydroxide hyper(name of metal)icoand hydroxide hiccup(name of metal)bear.

Stock

Of all the nomenclatures, this is the simplest. Here the name of the hydroxide is simply followed by the valence of the metal enclosed in parentheses and written in Roman numerals. Again for M(OH)5, for example, its stock nomenclature would be: (name of metal)(V) hydroxide. (V) then denotes (+5).

systematic

Finally, systematic nomenclature is characterized by resorting to multiplier prefixes (di-, tri-,…

Deja una respuesta

Tu dirección de correo electrónico no será publicada. Los campos obligatorios están marcados con *