Oxide Properties

What is the Oxide Properties: Let us know about Oxide Properties. Oxides They are a family of binary compounds where the interaction between the element and oxygen occurs. So an oxide has a very general formula of the EO type, where E is any element.

Depending on several factors, such as the electronic nature of E, its ionic radius and its valence, different types of oxides can be formed. Some are very simple, and others, such as Pb 3 O 4 , (called minium, arcazone or red lead); That is, they result from the combination of more than one simple oxide.

Oxide Properties
Oxide Properties

But the complexity of the oxide can go further. There are mixtures or structures in which more than one metal can interfere, and where the addition ratios are not stoichiometric. In the case of Pb 3 O 4 , the ratio of Pb / O is equal to 3/4, of which both the numerator and the denominator are integers.

Ratios in non-stoichiometric oxides are decimal numbers. E 0.75 is an example of O 1.78 , a hypothetical non-stoichiometric oxide. This phenomenon occurs with so-called metal oxides, especially with transition metals (Fe, Au, Ti, Mn, Zn, etc.).

However, there are oxides whose characteristics are very simple and different, as is the ionic or covalent character. The oxides where the ionic character predominates are composed of C + and the anion O 2 ; and those purely covalent, simple (E-O) or double (E=O) links.

What dictates the ionic character of an oxide is the electronegativity difference between E and O. When E is a very electronegative metal, Eo will have a high ionic character. Whereas if E is electronegative, then a non-metal, its EO oxide will be covalent.

This property defines many others exhibited by oxides, as is their ability to form bases or acids in aqueous solution. From here the so-called basic and acid oxides arise. Those that either do not behave, or show both characteristics, are neutral or amphoteric oxides.


There are three ways to refer to oxides (which also apply to many other compounds). These are true regardless of the ionic character of EO oxides, so their names say nothing about their properties or structures.

systematic naming

To know about Oxide Properties, now know systematic naming. Looking at the oxides EO, E 2 O, E 2 O 3 and EO 2 , at first glance you can’t know what’s behind your chemical formulas. However, the numbers reflect the stoichiometric ratio or E/O ratio. They can be named from these numbers, even if it is not specified with which valence the “e” works.

The number of atoms for both E and O is indicated by Greek number 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 systematic nomenclature are:

– This hair bale e (EO) oxide.

– This hairpin oxide di e (e) 2 o).

– Oxide of tri di e (e) 2 O 3 ).

– di e oxide (EO) 2 ).

Then applying this nomenclature to Pb 3 O 4 , the red oxide of the first image, we have:

Pb 3 O 4 : Oxide of tetra lead tri .

For many mixed oxides, or with high stoichiometric ratios, it is very useful to resort to systematic nomenclature to name them.

stock nomenclature


To know about Oxide Properties, now you know Valencia. Although it is not known which element is E, it is sufficient to know with the E/O ratio what valence it is using in its oxide. How? Through 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, O has a charge of -2 because it is O2- , and E must provide n+ so that it neutralizes the negative charges of the oxide ion.

For example, in EO the atom E acts with valence +2. Why? Because otherwise it could not neutralize E’s weight -2 only for O 2 Or, E has valence +1, because the charge +2 must be split between two atoms of E.

And in E 2 O 3 , the negative charges contributed by O must be calculated first. Since there are three of them, then: 3(-2) = -6. To neutralize the load -6 it is necessary that E provides +6, but because there are two of them, +6 is divided by two, leaving E with a wellness of +3.

mental rule

O always has valence -2 in the oxide (unless it is a peroxide or a superoxide). So the only rule to determine the validity of E is to take into account only the number that accompanies O. E, on the other hand, will be accompanied by the number 2, and if not, it means that there was a simplification.

For example, the value of E in EO is +1, because even if it is not written, there is only one O. And 2 for EO , ​​with 2 in the absence of E, was a simplification, and it must be multiplied by 2 to represent this. Thus, the formula remains from E. 2 O 4 and the value of E is +4.

However, this rule fails for some oxides such as Pb 3 O 4 . Therefore, it is always necessary to do neutrality calculations.

what does that involve?

To know about Oxide Properties, now know what does that involve? Once a hand has the validity of the e, in stock nomenclature it is specified within parentheses and with Roman numerals. Of all the nomenclatures, this is the simplest and most accurate in relation to the electronic properties of oxides.

If, on the other hand, E has only one valence (which can be found in the periodic table), then it is not specified.

Thus, for the oxide EO if E has valence +2 and +3, it is called: oxide OFF (name of E) (II). But if E has only valence +2, then its oxide is called: oxide (name of E).

traditional nomenclature

In order to refer to the name of the oxide, the suffix -co or -so, for major or minor valences, must be added to their Latin names. If there are more than two, prefix for smallest, and -at, largest for largest.

For example, Leeds valence works with +2 and +4. Its valence in PbO is +2, so it is called: plumb oxide. While P.B.O. 2 It is called: plumbico oxide.

and PB. 3 O 4 , how is it called according to the last two nomenclatures? It doesn’t have a name. Why? Because PB 3 O 4 actually consists of the mixture 2 [PbO] [PbO 2 ]; That is, the red solid has a double concentration of PbO.

For this reason it would be wrong to try to assign a name to Pb 3 O 4 that does not include systematic nomenclature or popular slang.

types of oxides

To know about Oxide Properties, now know the types of oxides. Depending on which part of the periodic table E and hence, its electronic nature, one type of oxide or the other can be made. From here a number of criteria arise to assign them a type of function, but the most important are those related to their acidity or basicity.

basic oxide

To know about Oxide Properties, now you know basic oxide. Basic oxides when dissolved in water produce a basic solution being ionic, metallic and more importantly. To experimentally determine if an oxide is basic, it must be added to a container with water and the universal indicator dissolved in it. It should be green in color, pH neutral before the oxide is added.

Once the oxide is added to water, if its color changes from green to blue, it means that the pH has become basic. This is because it establishes an equilibrium of solubility between the hydroxide formed and the water:

EO (S) + H 2 O (l) => E (OH) 2 (S) <=> A 2+ (AC) + OH  (AQ)

Although the oxide is insoluble in water, it is sufficient to a small extent to modify the pH. Some basic oxides are so soluble that they produce caustic hydroxides such as NaOH and KOH. That is, the oxides of sodium and potassium, Na 2 O and K 2 Or, they are very basic. Note the valency of +1 for both metals.

acid oxide

Acid oxides are characterized by being a non-metallic element, are covalent, and also produce acidic solutions with water. Again, its acidity can be checked with a universal indicator. If this time the oxide is mixed with water and its green color turns red, then it is an acid oxide.

What is the reaction? Following:

EO 2 (s) + H 2 O (L) => H 2 EO 3 (AQ)

An example of an acid oxide, which is not a solid, but a gas, is CO2 . When it dissolves in water, it forms carbonic acid:

CO 2 (g) + H 2 O (l) <=> H 2 CO 3 (aq)

Also, C.O. 2 It does not consist of ions or 2- and C 4+ , ​​but in a molecule formed by covalent bonds: O = C = O. This is probably one of the biggest differences between basic oxides and acids.

neutral oxide

These oxides do not turn water green at neutral pH; That is, they form neither hydroxide nor acid in aqueous solution. Some of them are: N 2 O, NO and CO. Like CO, they have covalent bonds that can be characterized by Lewis structures or any link theory.

amphoteric oxide

Another way to classify oxides is based on whether or not they react with an acid. Water is a very weak acid (and also a base), so amphoteric oxides do not exhibit “both sides”. 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 their reaction with acids or bases:

to 2 O 3 (s) + 3 H 2 sw 4 (ac) => Al 2 (so 4 ) 3 (ac) + 3 H 2 O (l)

To 2 O 3 (s) + 2NOH (AC) + 3 H 2 O (L) => 2 nal (OH) 4 (aq)

The Al 2 (SO 4 ) 3 is the aluminum sulfate salt, and NaAl (OH) 4 dissociates to a complex salt called sodium tetrahydroxyan.

Hydrogen oxide, H or (water), is also amphoteric, and is evidenced in its ionization equilibrium :

2 O (L) <=> H 3 O + (AC) + OH  (AQ)

mixed oxides

Mixed oxides are those that consist of a mixture of one or more oxides in a single solid. Pb 3 O 4 This is an example of them. Magnetite, believed to be 3 O 4 , is also another example of a mixed oxide. Astha 3 O 4 It is a mixture of FeO and Fe 2 O 3 in 1: 1 ratio (as opposed to Pb 3 O 4 ).

The mixture can be more complex, thus producing a rich variety of oxide minerals.


The properties of oxides depend on their type. Oxides can be ionic (E n+ O 2- ), such as CaO (Ca 2+ O 2- ), or covalent, as SO 2 , O = S = O.

From this fact, and the tendency of the elements to react with acids or bases, several properties are aggregated for each oxide.

In addition, the above physical properties such as melting and boiling points are reflected. Ionic oxides form crystalline structures that are very resistant to heat, so they have high melting points (over 1000 C), while covalent ones melt at low temperatures, or even gases or liquids.

How are they made?

contact with an atmosphere rich in oxygen, or requires heat (such as the flame of a cigarette lighter). That is, when an object burns, it reacts with oxygen (as long as it is present in the air).

If a piece of phosphorus is taken, for example, and placed in a flame, it burns and will form the corresponding oxide:

4 p (s) + 5 o 2 (g) => p 4 o 10 (s)

During this process some solids, such as calcium, may burn with a bright and colorful flame.

Another example is derived from wood or any organic material that contains carbon:

C (S) + O 2 (g) => CO 2 (G)

But if there is a lack of oxygen, instead of CO, CO 2 is formed :

C (s) + 1/2O 2 (g) => CO (g)

Note how the C/O ratio is used to describe the various oxides.

examples of oxides

The upper image corresponds to the covalent oxide structure I 2 O 5 , the most stable form of iodine. Note its simple and double bonds, as well as the formal charges of I and oxygen on its sides.

Halogen oxides are characterized by being covalent and very reactive, as are the cases of O 2 F 2 (F-O-O-F) and OF 2 (F-O-F). Chlorine dioxide, ClO , for example, is the only chlorine oxide that is synthesized on industrial scales.

Because the halogens form covalent oxides, their “imaginary” value is calculated through the principle of electroneutrality.

transition metal oxides

In addition to halogen oxides, we have oxides of transition metals:

-CoO: cobalt oxide (II); cobalt oxide; U cobalt monoxide.

-HgO: mercury oxide (II); mercuric oxide; u mercury monoxide.

-Ag 2 O: silver oxide; silver oxide; or Dipl monoxide.

-Au 2 O 3 : gold oxide (III); aureus oxide; or dirotrioxide.

additional examples

-B 2 O 3 : boron oxide; boric oxide; or dubro trioxide.

-chlorine 2 O 7 : chlorine oxide (VII); perchloric oxide; Dichloro heptoxide.

-NO: nitrogen oxides (II); nitric oxide; nitrogen monoxide.


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