The oxides They are a family of binary compounds where there are interactions between the element and oxygen. So an oxide has a very general formula of the type EO, where E is any element.
Depending on many factors, such as the electronic nature of E, its ionic radius, and its valences, various types of oxides can form. Some are very simple, and others, like Pb3OR4, (called minium, arcazón or red lead) are mixed; that is, they result from the combination of more than one simple oxide.
But the complexity of the oxides can go further. There are mixtures or structures in which more than one metal can intervene, and where also the proportions are not stoichiometric. In the case of Pb3OR4, the Pb / O ratio is equal to 3/4, of which both the numerator and denominator are whole numbers.
In non-stoichiometric oxides the proportions are decimal numbers. The E0.75OR1.78, is an example of a hypothetical non-stoichiometric oxide. This phenomenon occurs with the so-called metal oxides, especially with transition metals (Fe, Au, Ti, Mn, Zn, etc.).
However, there are oxides whose characteristics are much simpler and differentiable, such as ionic or covalent character. In those oxides where the ionic character predominates, they will be composed of cations E+ and anions Otwo-; and those purely covalent, the single bonds (E-O) or double bonds (E = O).
What dictates the ionic character of an oxide is the electronegativity difference between E and O. When E is a highly electropositive metal, then EO will have a high ionic character. Whereas if E is electronegative, namely a nonmetal, its oxide EO will be covalent.
This property defines many others exhibited by oxides, such as their ability to form bases or acids in aqueous solution. From here arise the so-called basic and acid oxides. Those that do not behave like either of the two, or that on the contrary show both characteristics, are neutral or amphoteric oxides.
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There are three ways to name oxides (which apply to many other compounds as well). These are correct regardless of the ionic character of the EO oxide, so their names say nothing about its properties or structures.
Given the oxides EO, EtwoO, EtwoOR3 and EOtwo, at first glance it is not possible to know what is behind their chemical formulas. However, the numbers indicate the stoichiometric ratios or the E / O ratio. From these numbers they can be given names although it is not specified with what valence it "works" E.
The numbers of atoms for both E and O are denoted by the Greek numbering prefixes. In this way, mono- means that there is only one atom; di-, two atoms; tri-, three atoms, and so on.
So, the names of the previous oxides according to the systematic nomenclature are:
-MonkeyE oxide (EO).
-Monkeyoxide of gaveE (EtwoOR).
-Trioxide of gaveE (EtwoOR3).
-GaveE oxide (EOtwo).
Applying then this nomenclature for Pb3OR4, the red oxide of the first image, we have
Pb3OR4: tetraoxide of trilead.
For many mixed oxides, or with high stoichiometric ratios, it is very useful to use the systematic nomenclature to name them.
Although it is not known which element is E, it is enough with the E / O ratio to know what valence you are using in your oxide. How? By the principle of electroneutrality. This requires that the sum of the charges of the ions in a compound must be equal to zero..
This is done by assuming a high ionic character for any oxide. Thus, the O has charge -2 because it is the Otwo-, and E must contribute n + so that it neutralizes the negative charges of the oxide anion.
For example, in EO the E atom works with valence +2. Why? Because otherwise it could not neutralize the charge -2 of the only O. For the EtwoOr, E has valence +1, since the +2 charge must be divided between the two E atoms..
And in the EtwoOR3, The negative charges contributed by O must be calculated first. Since there are three of them, then: 3 (-2) = -6. To neutralize the -6 charge, the E's are required to contribute +6, but since there are two of them, +6 is divided by two, leaving E with a valence of +3.
O always has a -2 valence in oxides (unless it is a peroxide or superoxide). So a mnemonic rule to determine the valence of E is simply to take into account the number that accompanies the O. E, on the other hand, will have the number 2 accompanying it, and if not, it means that there was a simplification.
For example, in EO the valence of E is +1, because even if it is not written, there is only one O. And for EOtwo, As there is no 2 accompanying E, there was a simplification, and for it to appear it must be multiplied by 2. Thus, the formula remains as EtwoOR4 and the valence of E is then +4.
However, this rule fails for some oxides, such as Pb3OR4. Therefore, it is always necessary to perform neutrality calculations.
Once the valence of E is at hand, the stock nomenclature consists of specifying it within parentheses and with Roman numerals. Of all the nomenclatures this is the simplest and most accurate with respect to the electronic properties of oxides..
If E, on the other hand, has only one valence (which can be found in the periodic table), then it is not specified..
Thus, for oxide EO if E has valence +2 and +3, it is called: (name of E) (II) oxide. But if E only has valence +2, then its oxide is called: oxide of (name of E).
To mention the name of the oxides, the suffixes -ico or -oso must be added to their Latin names, for the larger or smaller valences. In the event that there are more than two, the prefixes -hipo, for the smallest, and -per, for the largest of all, are used.
For example, lead works with valences +2 and +4. In PbO it has a valence of +2, so it is called: plumbous oxide. While the PbOtwo it's called: lead oxide.
And the Pb3OR4, what is it called according to the two previous nomenclatures? It has no name. Why? Because the Pb3OR4 actually consists of a mixture 2 [PbO] [PbOtwo]; that is, the red solid has a double concentration of PbO.
For this reason it would be wrong to try to name the Pb3OR4 that does not consist of systematic nomenclature or popular slang.
Depending on which part of the periodic table E is located and, therefore, its electronic nature, one type of oxide or another can be formed. From this multiple criteria arise to assign them a type, but the most important are those related to their acidity or basicity..
Basic oxides are characterized by being ionic, metallic, and more importantly, generating a basic solution by dissolving in water. To determine experimentally whether an oxide is basic, it must be added to a container with water and a universal indicator dissolved in it. Its coloration before adding the oxide must be green, pH neutral.
Once the oxide has been added to the water, if its color changes from green to blue, it means that the pH has become basic. This is because it establishes a solubility balance between the hydroxide formed and the water:
EO (s) + HtwoO (l) => E (OH)two(s) <=> ANDtwo+(ac) + OH-(ac)
Although the oxide is insoluble in water, only a small portion dissolves to change the pH. Some basic oxides are so soluble that they generate caustic hydroxides such as NaOH and KOH. That is, the oxides of sodium and potassium, NatwoO and KtwoOr, they are very basic. Note the valence of +1 for both metals.
Acidic oxides are characterized by having a non-metallic element, they are covalent, and in addition, they generate acidic solutions with water. Again, its acidity can be checked with the universal indicator. If this time when adding the oxide to the water, its green color turns reddish, then it is an acid oxide.
What reaction takes place? The next:
EOtwo(s) + HtwoO (l) => HtwoEO3(ac)
An example of an acid oxide, which is not a solid, but a gas, is COtwo. When it dissolves in water, it forms carbonic acid:
COtwo(g) + HtwoO (l) <=> HtwoCO3(ac)
Also, the COtwo does not consist of anions Otwo- and C cations4+, but in a molecule formed by covalent bonds: O = C = O. This is perhaps one of the biggest differences between basic oxides and acids.
These oxides do not change the green color of water at neutral pH; that is, they do not form hydroxides or acids in aqueous solution. Some of them are: NtwoOR, NO and CO. Like CO, they have covalent bonds that can be illustrated by Lewis structures or any theory of bonding..
Another way to classify oxides depends on whether or not they react with an acid. Water is a very weak acid (and a base too), so amphoteric oxides do not exhibit "their two faces." These oxides are characterized by reacting with both acids and bases.
Aluminum oxide, for example, is an amphoteric oxide. The following two chemical equations represent its reaction with acids or bases:
To thetwoOR3(s) + 3HtwoSW4(ac) => Altwo(SW4)3(aq) + 3HtwoO (l)
To thetwoOR3(s) + 2NaOH (aq) + 3HtwoO (l) => 2NaAl (OH)4(ac)
The Altwo(SW4)3 is the aluminum sulfate salt, and the NaAl (OH)4 a complex salt called sodium tetrahydroxy aluminate.
Hydrogen oxide, HtwoO (water), is also amphoteric, and this is evidenced by its ionization balance:
HtwoO (l) <=> H3OR+(ac) + OH-(ac)
Mixed oxides are those that consist of the mixture of one or more oxides in the same solid. The Pb3OR4 is an example of them. Magnetite, Fe3OR4, it is also another example of a mixed oxide. Faith3OR4 is a mixture of FeO and FetwoOR3 in 1: 1 proportions (unlike Pb3OR4).
Mixtures can be more complex, thus creating a rich variety of oxide minerals.
The properties of oxides depend on their type. Oxides can be ionic (En+ORtwo-), such as CaO (Catwo+ORtwo-), or covalent, such as SOtwo, O = S = O.
From this fact, and from the tendency of elements to react with acids or bases, a number of properties are collected for each oxide..
Also, the above is reflected in the physical properties such as melting and boiling points. Ionic oxides tend to form crystalline structures that are very resistant to heat, so their melting points are high (above 1000ºC), while covalents melt at low temperatures, or they are even gases or liquids..
Oxides are formed when elements react with oxygen. This reaction can occur with simple contact with oxygen-rich atmospheres, or requires heat (such as a lighter flame). That is, when burning an object it reacts with oxygen (as long as it is present in the air).
If you take a piece of phosphorus, for example, and place it in the flame, it will burn and form the corresponding oxide:
4P (s) + 5Otwo(g) => P4OR10(s)
During this process some solids, such as calcium, can burn with a bright, colorful flame..
Another example is obtained by burning wood or any organic substance, which have carbon:
C (s) + Otwo(g) => COtwo(g)
But if there is insufficient oxygen, CO is formed instead of COtwo:
C (s) + 1 / 2Otwo(g) => CO (g)
Note how the C / O ratio serves to describe different oxides.
The upper image corresponds to the structure of covalent oxide ItwoOR5, the most stable form of iodine. Note their single and double bonds, as well as the formal charges of I and oxygens on their sides..
The oxides of halogens are characterized by being covalent and highly reactive, as such are the cases of OtwoFtwo (F-O-O-F) and OFtwo (F-O-F). Chlorine dioxide, ClOtwo, for example, it is the only chlorine oxide that is synthesized on an industrial scale.
Since halogens form covalent oxides, their "hypothetical" valences are calculated in the same way through the principle of electroneutrality..
In addition to the halogen oxides, there are the transition metal oxides:
-CoO: cobalt (II) oxide; cobalt oxide; u cobalt monoxide.
-HgO: mercury (II) oxide; mercuric oxide; u mercury monoxide.
-AgtwoO: silver oxide; silver oxide; or diplate monoxide.
-AutwoOR3: gold (III) oxide; auric oxide; or dioro trioxide.
-BtwoOR3: boron oxide; boric oxide; or diboron trioxide.
-CltwoOR7: chlorine oxide (VII); perchloric oxide; dichloro heptoxide.
-NO: nitrogen oxide (II); Nitric oxide; nitrogen monoxide.
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