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 cation M+ and the OH anion-, or as the covalent bond through the M-OH bond (lower image). In the first, the ionic bond occurs, while in the second, the covalent. This fact depends essentially 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.
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There are two main synthetic routes: by reacting the corresponding oxide with water, or with a strong base in an acidic medium:
MO + HtwoO => M (OH)two
MO + H+ + Oh- => M (OH)two
Only those metal oxides soluble in water react directly to form the hydroxide (first chemical equation). Others are insoluble and require acidic species to release M+, which then interacts with the OH- from strong bases (second chemical equation).
However, these strong bases are metal hydroxides NaOH, KOH and others from the group of alkali metals (LiOH, RbOH, CsOH). These are ionic compounds highly soluble in water, hence their OH- are free to participate in chemical reactions.
On the other hand, there are metal hydroxides that are insoluble and consequently are very weak bases. Some of them are even acidic, as is the case with telluric acid, Te (OH)6.
The hydroxide establishes a solubility equilibrium with the surrounding solvent. If it is water, for example, then the equilibrium is expressed as follows:
M (OH)two <=> Mtwo+(ac) + OH-(ac)
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 NaOH..
From the above it can be deduced that hydroxides exhibit very different properties, linked to the chemical structure and the interactions between metal and OH. Thus, although many are ionic, with varied crystalline structures, others have complex and disordered polymeric structures..
The hydroxyl ion is an oxygen atom covalently bonded to hydrogen. Thus, this can easily be represented as OH-. The negative charge is located on the oxygen, making this anion an electron donor species: a base.
If the OH- donates its electrons to a hydrogen, a molecule of H is formedtwoO. You can also donate your electrons to positively charged species: such as metal centers M+. Thus, a coordination complex is formed through the dative M-OH bond (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 marked ionic character (M+ Oh-). As the hydroxyl ion is the same in all hydroxides, the difference between all of them then lies in the cation that accompanies it.
Likewise, because this cation can come from any metal on the periodic table (groups 1, 2, 13, 14, 15, 16, or from the transition metals), the properties of such hydroxides vary enormously, although all contemplate in common some aspects.
In hydroxides, although they have coordination bonds, they have a latent ionic character. In some, such as NaOH, their ions are part of a crystal lattice made up of Na cations+ and OH anions- in proportions 1: 1; that is, for each Na ion+ there is an OH ion- counterpart.
Depending on the charge of the metal, there will be more or less OH anions- around it. For example, for a metal cation Mtwo+ there will be two OH ions- interacting with him: M (OH)two, what is sketched as HO- Mtwo+ Oh-. In the same way it happens with the metals M3+ 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 at work within the crystal lattice. Likewise, when hydroxides dissolve or melt, they can conduct electric current due to the mobility of their ions..
However, not all hydroxides have the same crystal lattices. Those with the most stable ones 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-, more soluble will be the same.
This explains why the solubility of alkali metal hydroxides increases as one moves down the group. Thus, the increasing order of solubilities in water for these is as follows: LiOH The OH- is a small anion, and as the cation becomes more voluminous, the crystal lattice energetically weakens. On the other hand, alkaline earth metals form less soluble hydroxides due to their higher positive charges. This is because the Mtwo+ attracts OH more strongly- compared to 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).two. The same reasoning can be applied for 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 the 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)two. He betwo+ it is a very small cation and its divalent charge makes it electrically very dense. Hydroxides M (OH)two react with acids to form an aqueous complex, 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)two + 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 the M ions+ dissolved, for example, Al3+, the corresponding hydroxide will precipitate: To the3+(aq) + 3OH-(ac) => Al (OH)3(s) But having an excess of OH- the hydroxide continues to react: Al (OH)3(s) + OH- => Al (OH)4-(ac) As a result, the new negatively charged complex is solvated by the surrounding water molecules, dissolving the white aluminum hydroxide solid. Those hydroxides that remain unchanged with the extra addition of base do not behave as acids and, therefore, are not amphoteric. Hydroxides can have crystalline structures similar to those of many salts or oxides; some simple, and others very complex. In addition, those where there is a decrease in ionic character may present metal 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 into account the coordination chemistry. Thus, each cation M+ it can be coordinated to a limited number of species. The bulkier it is, the greater the number of molecules of water or OH.- linked to it. From this arises the famous octahedron of coordination of many metals dissolved in water (or in any other solvent): M (OHtwo)6+n, where n is equal to the positive charge of the metal. Cr (OH)3, for example, it actually forms an octahedron. How? Considering the compound as [Cr (OHtwo)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 of negative charge and octahedral structure [Cr (OH)6]3-. The -3 charge is the result of the six negative charges of the OH-. Hydroxides can be considered as "hydrated oxides". However, in them "water" is in direct contact with M+; while in the hydrated oxides MO nHtwoOr, the water molecules are part of an external coordination sphere (they are not close to the metal). These water molecules can be extracted by heating a sample of hydroxide: M (OH)two + Q (heat) => MO + HtwoOR MO is the metal oxide formed as a result of the dehydration of the hydroxide. An example of this reaction is the one observed when cupric hydroxide, Cu (OH), is dehydratedtwo: Cu (OH)two (blue) + Q => CuO (black) + HtwoOR What is the proper way to mention hydroxides? The IUPAC proposed 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 convenient or practical to mention it in one way or another. The traditional nomenclature is simply to add the suffix -ico to the highest valence of the metal; and the suffix -oso to the lowest. Thus, for example, if the metal M has valences +3 and +1, the hydroxide M (OH)3 it will be called hydroxide (metal name)ico, while MOH hydroxide (metal name)bear. To determine the valence 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, it also has +4 and +2, then the names of its hydroxides with the highest and lowest valences are: hydroxide hyper(metal name)ico, and hydroxide hiccup(metal name)bear. 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 the M (OH)5, for example, your stock nomenclature would be: (metal name) hydroxide (V). (V) denotes then (+5). Finally, the systematic nomenclature is characterized by resorting to multiplying prefixes (di-, tri-, tetra-, penta-, hexa-, etc.). These prefixes are used to specify both the number of metal atoms and OH ions.-. In this way, the M (OH)5 named as: (metal name) pentahydroxide. In the case of Hgtwo(OH)two, for example, it would be dimercuric dihydroxide; one of the hydroxides whose chemical structure is complex at first glance. Some examples of hydroxides and their corresponding nomenclatures are as follows: -NaOH (Sodium Hydroxide) -Ca (OH) 2 (Calcium hydroxide) -Fe (OH)3. (Ferric hydroxide; iron (III) hydroxide; or iron trihydroxide) -V (OH)5 (Pervanadic hydroxide; Vanadium (V) hydroxide; or vanadium pentahydroxide). -Sn (OH)4 (Stannic hydroxide; tin (IV) hydroxide; or tin tetrahydroxide). -Ba (OH)two (Barium hydroxide or barium dihydroxide). -Mn (OH)6 (Manganic hydroxide, manganese (VI) hydroxide or manganese hexahydroxide). -AgOH (Silver hydroxide, silver hydroxide or silver hydroxide). Note that for this compound there is no distinction between the stock and systematic nomenclatures. -Pb (OH)4 (Lead hydroxide, lead (IV) hydroxide, or lead tetrahydroxide). -LiOP (Lithium Hydroxide). -Cd (OH) 2 (Cadmium Hydroxide) -Ba (OH)two (Barium hydroxide) -Chromium hydroxideAmphotericism
Structures
Dehydration reaction
Nomenclature
Traditional
Stock
Systematic
Examples of hydroxides
References
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