- Properties of non-metallic oxides
- Basicity
- Amphotericism
- Nomenclature
- Traditional nomenclature
- Systematic nomenclature
- Stock nomenclature
- Calculation of the valence number
- How are they formed?
- Direct reaction of metal with oxygen
- Reaction of metal salts with oxygen
- Applications
- Examples
- Iron oxides
- Alkali and alkaline earth oxides
- Group IIIA oxides (13)
- References
Metal oxides are inorganic compounds made up of metal cations and oxygen. They generally comprise a vast number of ionic solids, in which the oxide anion (O 2–) interacts electrostatically with the M + species.
M + is like this any cation that derives from the pure metal: from the alkali and transition metals, with the exception of some noble metals (such as gold, platinum and palladium), to the heaviest elements of the p block of the table periodic (such as lead and bismuth).
Source: Pixabay.
The image above shows an iron surface covered by reddish crusts. These "scabs" are what is known as rust or rust, which in turn represent visual evidence of the oxidation of the metal as a result of the conditions of its environment. Chemically, rust is a hydrated mixture of iron (III) oxides.
Why does the oxidation of the metal lead to the degradation of its surface? This is due to the incorporation of oxygen within the crystal structure of the metal.
When this happens, the volume of the metal increases and the original interactions weaken, causing the solid to rupture. Likewise, these cracks allow more oxygen molecules to penetrate the internal metal layers, completely eating away at the piece from the inside.
However, this process occurs at different speeds and depends on the nature of the metal (its reactivity) and the physical conditions that surround it. Therefore, there are factors that accelerate or slow down the oxidation of the metal; two of them are the presence of humidity and pH.
Why? Because the oxidation of metal to produce a metal oxide involves a transfer of electrons. These “travel” from one chemical species to another as long as the environment facilitates it, either by the presence of ions (H +, Na +, Mg 2+, Cl -, etc.), which modify the pH, or by the water molecules that provide the means of transport.
Analytically, the tendency of a metal to form the corresponding oxide is reflected in its reduction potentials, which reveal which metal reacts faster compared to another.
Gold, for example, has a much greater reduction potential than iron, which is why it shines with its characteristic golden glow without an oxide to dull it.
Properties of non-metallic oxides
Magnesium oxide, a metal oxide.
The properties of metal oxides vary according to the metal and how it interacts with the O 2– anion. This means that some oxides have higher densities or solubilities in water than others. However, they all have in common the metallic character, which is inevitably reflected in their basicity.
In other words: they are also known as basic anhydrides or basic oxides.
Basicity
The basicity of metal oxides can be verified experimentally by using an acid-base indicator. How? Adding a small piece of the oxide to an aqueous solution with some dissolved indicator; This can be the liquefied juice of the purple cabbage.
Having then the range of colors depending on the pH, the oxide will turn the juice to bluish colors, corresponding to basic pH (with values between 8 and 10). This is due to the fact that the dissolved portion of the oxide releases OH - ions into the medium, these being responsible for the change in pH in said experiment.
Thus, for an oxide MO that is solubilized in water, it is transformed into metallic hydroxide (a “hydrated oxide”) according to the following chemical equations:
MO + H 2 O => M (OH) 2
M (OH) 2 <=> M 2+ + 2OH -
The second equation is the solubility equilibrium of the hydroxide M (OH) 2. Note that the metal has a 2+ charge, which also means that its valence is +2. The valence of the metal is directly related to its tendency to gain electrons.
In this way, the more positive the valence, the greater its acidity. In the case that M had a valence of +7, then the oxide M 2 O 7 would be acidic and not basic.
Amphotericism
Metal oxides are basic, however they do not all have the same metallic character. How do you know? Locating the metal M on the periodic table. The further you are to the left of it, and in low periods, the more metallic it will be and therefore the more basic your oxide will be.
At the border between basic and acidic oxides (non-metallic oxides) are amphoteric oxides. Here the word 'amphoteric' means that the oxide acts both as a base and as an acid, which is the same as in aqueous solution it can form the hydroxide or the aqueous complex M (OH 2) 6 2+.
The aqueous complex is nothing more than the coordination of n water molecules with the metal center M. For the M (OH 2) 6 2+ complex, the metal M 2+ is surrounded by six water molecules, and can be considered as a hydrated cation. Many of these complexes show intense colorations, such as those observed for copper and cobalt.
Nomenclature
How are metal oxides named? There are three ways to do it: traditional, systematic, and stock.
Traditional nomenclature
To correctly name the metal oxide according to the rules governed by the IUPAC, it is necessary to know the possible valences of the metal M. The largest (the most positive) is assigned the suffix -ico to the name of the metal, while the minor, the prefix –oso.
Example: given the +2 and +4 valences of the metal M, its corresponding oxides are MO and MO 2. If M off the lead, Pb, then oxide PbO Plumb would bear, and PbO 2 oxide PLUMB ico. If the metal has only one valence, its oxide is named with the suffix –ico. Thus, Na 2 O is sodium oxide.
On the other hand, the prefixes hypo- and per- are added when there are three or four valences available for the metal. Thus, the Mn 2 O 7 is oxide per Mangan ico, because Mn has +7 valence, most of all.
However, this type of nomenclature presents certain difficulties and is usually the least used.
Systematic nomenclature
In it, the number of M and oxygen atoms that make up the chemical formula of the oxide are considered. From them, it is assigned the corresponding prefixes mono-, di-, tri-, tetra-, etc.
Taking the three recent metal oxides as an example, PbO is lead monoxide; PbO 2 lead dioxide; and Na 2 O is disodium monoxide. In the case of rust, Fe 2 O 3, its respective name is di iron trioxide.
Stock nomenclature
Unlike the other two nomenclatures, in this one, the valence of the metal is more important. Valence is specified by Roman numerals in parentheses: (I), (II), (III), (IV), etc. The metal oxide is then named metal (n) oxide.
Applying the stock nomenclature for the previous examples, we have:
-PbO: lead (II) oxide.
-PbO 2: lead (IV) oxide.
-Na 2 O: sodium oxide. Since it has a unique valence of +1, it is not specified.
-Fe 2 O 3: iron (III) oxide.
-Mn 2 O 7: manganese (VII) oxide.
Calculation of the valence number
But, if you don't have a periodic table with the valences, how can you determine them? For this we must remember that the anion O 2– contributes two negative charges to the metal oxide. Following the principle of neutrality, these negative charges must be neutralized with the positive ones of the metal.
Therefore, if the number of oxygens is known from the chemical formula, the valence of the metal can be determined algebraically so that the sum of the charges is zero.
Mn 2 O 7 has seven oxygens, so its negative charges are equal to 7x (-2) = -14. To neutralize the negative charge of -14, manganese must contribute +14 (14-14 = 0). Posing the mathematical equation we have then:
2X - 14 = 0
The 2 comes from the fact that there are two manganese atoms. Solving and solving for X, the valence of the metal:
X = 14/2 = 7
In other words, each Mn has a valence of +7.
How are they formed?
Moisture and pH directly influence the oxidation of metals into their corresponding oxides. The presence of CO 2, acidic oxide, can dissolve sufficiently in the water that covers the metal part to accelerate the incorporation of oxygen in anionic form to the crystal structure of the metal.
This reaction can also be accelerated with an increase in temperature, especially when it is desired to obtain the oxide in a short time.
Direct reaction of metal with oxygen
Metal oxides are formed as a product of the reaction between the metal and the surrounding oxygen. This can be represented by the chemical equation below:
2M (s) + O 2 (g) => 2MO (s)
This reaction is slow, since oxygen has a strong O = O double bond and the electronic transfer between it and the metal is inefficient.
However, it accelerates considerably with an increase in temperature and surface area. This is due to the fact that the necessary energy is provided to break the O = O double bond, and as there is a greater area, the oxygen moves uniformly throughout the metal, colliding at the same time with the metal atoms.
The greater the amount of reacting oxygen, the greater the resulting valence or oxidation number for the metal. Why? Because oxygen takes more and more electrons from the metal, until it reaches the highest oxidation number.
This can be seen for copper, for example. When a piece of metallic copper reacts with a limited amount of oxygen, Cu 2 O is formed (copper (I) oxide, cuprous oxide, or dicobre monoxide):
4Cu (s) + O 2 (g) + Q (heat) => 2Cu 2 O (s) (red solid)
But when it reacts in equivalent amounts, CuO (copper (II) oxide, cupric oxide or copper monoxide) is obtained:
2Cu (s) + O 2 (g) + Q (heat) => 2CuO (s) (black solid)
Reaction of metal salts with oxygen
Metal oxides can be formed through thermal decomposition. For this to be possible, one or two small molecules must be released from the starting compound (a salt or a hydroxide):
M (OH) 2 + Q => MO + H 2 O
OLS 3 + Q => MO + CO 2
2M (NO 3) 2 + Q => MO + 4NO 2 + O 2
Note that H 2 O, CO 2, NO 2 and O 2 are the released molecules.
Applications
Due to the rich composition of metals in the earth's crust, and the oxygen in the atmosphere, metal oxides are found in many mineralogical sources, from which a solid basis for the manufacture of new materials can be obtained.
Each metal oxide finds very specific uses, from nutritional (ZnO and MgO) to as cement additives (CaO), or simply as inorganic pigments (Cr 2 O 3).
Some oxides are so dense that controlled layer growth can protect an alloy or metal from further oxidation. Studies have even revealed that the oxidation of the protective layer continues as if it were a liquid that covers all the cracks or superficial defects of the metal.
Metal oxides can take on fascinating structures, either as nanoparticles or as large polymer aggregates.
This fact makes them the object of studies for the synthesis of smart materials, due to their large surface area, which is used to design devices that respond to the least physical stimulus.
In addition, metal oxides are the raw material for many technological applications, from mirrors and ceramics with unique properties for electronic equipment, to solar panels.
Examples
Iron oxides
2Fe (s) + O 2 (g) => 2FeO (s) iron (II) oxide.
6FeO (s) + O 2 (g) => 2Fe 3 O 4 (s) magnetic iron oxide.
Fe 3 O 4, also known as magnetite, is a mixed oxide; This means that it consists of a solid mixture of FeO and Fe 2 O 3.
4Fe 3 O 4 (s) + O 2 (g) => 6Fe 2 O 3 (s) iron (III) oxide.
Alkali and alkaline earth oxides
Both alkali and alkaline earth metals have only one oxidation number, so their oxides are more “simple”:
-Na 2 O: sodium oxide.
-Li 2 O: lithium oxide.
-K 2 O: potassium oxide.
-CaO: calcium oxide.
-MgO: magnesium oxide.
-BeO: beryllium oxide (which is an amphoteric oxide)
Group IIIA oxides (13)
Group IIIA elements (13) can form oxides only with an oxidation number of +3. Thus, they have the chemical formula M 2 O 3 and their oxides are the following:
-Al 2 O 3: aluminum oxide.
-Ga 2 O 3: gallium oxide.
-In 2 O 3: indium oxide.
And finally
-Tl 2 O 3: thallium oxide.
References
- Whitten, Davis, Peck & Stanley. Chemistry. (8th ed.). CENGAGE Learning, p 237.
- AlonsoFormula. Metal Oxides. Taken from: alonsoformula.com
- Regents of the University of Minnesota. (2018). Acid-base Characteristics of Metal and Nonmetal Oxides. Taken from: chem.umn.edu
- David L. Chandler. (April 3, 2018). Self-healing metal oxides could protect against corrosion. Taken from: news.mit.edu
- The Physical States and Structures of Oxides. Taken from: wou.edu
- Quimitube. (2012). Iron oxidation. Taken from: quimitube.com
- Chemistry LibreTexts. Oxides. Taken from: chem.libretexts.org
- Kumar M. (2016) Metal Oxide Nanostructures: Growth and Applications. In: Husain M., Khan Z. (eds) Advances in Nanomaterials. Advanced Structured Materials, vol 79. Springer, New Delhi