CHROMIUM: PROPERTIES, CHARACTERISTICS AND USES - CHEMISTRY - 2025

The chromium (Cr) is a metal element of Group 6 (VIB) of the periodic table. Tons of this metal are produced annually through its extraction from the chromite mineral of iron or magnesium (FeCr ₂ O ₄, MgCr ₂ O ₄), which are reduced with coal to obtain the metal. It is very reactive, and only under very reducing conditions is it in its pure form.

Its name derives from the Greek word 'chroma', which means color. It was given this name because of the multiple and intense colors exhibited by chromium compounds, whether inorganic or organic; from black solids or solutions to yellow, orange, green, violet, blue and red.

Chrome crocodile. Silver Crocodile Chrome Metal Model Alligator. Source: Maxpixel

However, the color of metallic chrome and its carbides are silvery grayish. This characteristic is exploited in the chrome plating technique to give many structures silver sparkles (like the ones seen in the crocodile in the image above). Thus, by “bathing with chrome” the pieces are given luster and great resistance to corrosion.

Chromium in solution reacts rapidly with oxygen in the air to form oxides. Depending on the pH and the oxidative conditions of the medium, it can acquire different oxidation numbers, with (III) (Cr ³⁺) being the most stable of all. Consequently, the green chromium (III) oxide (Cr ₂ O ₃) is the most stable of its oxides.

These oxides can interact with other metals in the environment, causing, for example, the pigment Siberian red lead (PbCrO ₄). This pigment is yellow-orange or red (according to its alkalinity), and from it the French scientist Louis Nicolas Vauquelin isolated metallic copper, which is why he is awarded as its discoverer.

Its minerals and oxides, as well as a tiny portion of metallic copper make this element rank 22 of the most abundant in the earth's crust.

The chemistry of chromium is very diverse because it can form bonds with almost the entire entire periodic table. Each of its compounds exhibits colors that depend on the oxidation number, as well as the species that interact with it. Likewise, it forms bonds with carbon, intervening in a large number of organometallic compounds.

Characteristics and properties

Chromium is a silver metal in its pure form, with an atomic number of 24 and a molecular weight of about 52 g / mol (⁵² Cr, its most stable isotope).

Given its strong metallic bonds, it has high melting points (1907 ºC) and boiling points (2671 ºC). Also, its crystalline structure makes it a very dense metal (7.19 g / mL).

It does not react with water to form hydroxides, but it does react with acids. It oxidizes with the oxygen in the air, generally producing chromic oxide, which is a widely used green pigment.

These oxide layers create what is known as passivation, protecting the metal from further corrosion, since oxygen cannot penetrate the metal sinus.

Its electronic configuration is 4s ¹ 3d ⁵, with all electrons unpaired, and therefore, it exhibits paramagnetic properties. However, the mating of electronic spins can occur if the metal is subjected to low temperatures, acquiring other properties such as antiferromagnetism.

Chromium chemical structure

By Original PNGs by Daniel Mayer, DrBob, traced in Inkscape by User: Stannered (Crystal stucture), via Wikimedia Commons

What is the structure of chrome metal? In its pure form, chromium assumes a body-centered cubic crystal structure (cc or bcc). This means that the chromium atom is located in the center of a cube, whose edges are occupied by other chromes (as in the image above).

This structure is responsible for chromium having high melting and boiling points, as well as high hardness. Copper atoms overlap their s and d orbitals to form conduction bands according to band theory.

Thus, both bands are half full. Why? Because its electronic configuration is 4s ¹ 3d ⁵ and as the s orbital it can house two electrons, and the d orbitals ten. Then only half of the bands formed by their overlaps are occupied by electrons.

With these two perspectives - the crystalline structure and the metallic bond - many of the physical properties of this metal can be explained in theory. However, neither explains why chromium can have various oxidation states or numbers.

This would require a deep understanding of the stability of the atom with respect to electronic spins.

Oxidation number

Since the electron configuration of chromium is 4s ¹ 3d ^{5 it} can gain up to one or two electrons (Cr ^1– and Cr ^2–), or lose them to acquire different oxidation numbers.

Thus, if chromium loses an electron, it would be 4s ⁰ 3d ⁵; if he loses three, 4s ⁰ 3d ³; and if it loses all of them, or what is the same, it would be isoelectronic to argon.

Chromium does not lose or gain electrons by mere whim: there has to be a species that donates or accepts them in order to go from one oxidation number to another.

Chromium has the following oxidation numbers: -2, -1, 0, +1, +2, +3, +4, +5, and +6. Of these, +3, Cr ³⁺, is the most stable and therefore predominant of all; followed by +6, Cr ⁶⁺.

Cr (-2, -1 and 0)

Chromium is highly unlikely to gain electrons, since it is a metal, and therefore its nature is to donate them. However, it can coordinate with ligands, that is, molecules that interact with the metal center through a dative bond.

One of the best known is carbon monoxide (CO), which forms the hexacarbonyl compound of chromium.

This compound has the molecular formula Cr (CO) ₆, and since the ligands are neutral and do not provide any charge, then Cr has an oxidation number of 0.

This can also be observed in other organometallic compounds such as bis (benzene) chromium. In the latter, chromium is surrounded by two benzene rings in a sandwich-like molecular structure:

By Ben Mills, from Wikimedia Commons

Many other Cr (0) compounds can arise from these two organometallic compounds.

Salts have been found where they interact with sodium cations, which implies that Cr must have a negative oxidation number to attract positive charges: Cr (-2), Na ₂ and Cr (-1), Na ₂.

Cr (I) and Cr (II)

Cr (I) or Cr ¹⁺ is produced by the oxidation of the organometallic compounds just described. This is achieved by oxidizing ligands, such as CN or NO, thus forming, for example, compound K ₃.

Here the fact that there are three K ⁺ cations implies that the chromium complex has three negative charges; likewise the ligand CN ^- contributes five negative charges, so that between Cr and NO they must add two positive charges (-5 + 2 = -3).

If NO is neutral, then it is Cr (II), but if it has a positive charge (NO ⁺), then it is Cr (I).

On the other hand, the Cr (II) compounds are more abundant, among them the following: chromium (II) chloride (CrCl ₂), chromous acetate (Cr ₂ (O ₂ CCH ₃) ₄), chromium oxide (II) (CrO), chromium (II) sulfide (CrS), and more.

Cr (III)

Of all, it is the one with the greatest stability, since it is in fact the product of many oxidative reactions of chromate ions. Perhaps its stability is due to its d ³ electronic configuration, in which three electrons occupy three lower-energy d orbitals compared to the other two more energetic ones (d-orbital doubling).

The most representative compound of this oxidation number is chromium (III) oxide (Cr ₂ O ₃). Depending on the ligands that coordinate to it, the complex will exhibit one color or another. Examples of these compounds are: Cl, Cr (OH) ₃, CrF ₃, ³⁺, etc.

Although the chemical formula does not show it at first glance, chromium usually has an octahedral coordination sphere in its complexes; that is, it is located in the center of an octahedron where its vertices are positioned by the ligands (six in total).

Cr (IV) and Cr (V)

The compounds in which Cr ⁵⁺ participates are very few, due to the electronic instability of said atom, in addition to the fact that it is easily oxidized to Cr ⁶⁺, much more stable as it is isoelectronic with respect to argon noble gas.

However, Cr (V) compounds can be synthesized under certain conditions, such as high pressure. Likewise, they tend to decompose at moderate temperatures, which makes their possible applications impossible as they do not have thermal resistance. Some of them are: CrF ₅ and K ₃ (O ₂ ^2- is the peroxide anion).

On the other hand, Cr ⁴⁺ is relatively more stable, being able to synthesize its halogenated compounds: CrF ₄, CrCl ₄ and CrBr ₄. However, they are also susceptible to being decomposed by redox reactions to produce chromium atoms with better oxidation numbers (such as +3 or +6).

Cr (VI): the chromate-dichromate pair

2 ^2- + 2H ⁺ (Yellow) => ^2- + H ₂ O (Orange)

The above equation corresponds to the acid dimerization of two chromate ions to produce dichromate. The variation in pH causes a change in the interactions around the metallic center of Cr ⁶⁺, also evident in the color of the solution (from yellow to orange or vice versa). Dichromate consists of an O ₃ Cr-O-CrO ₃ bridge.

Cr (VI) compounds have the characteristics of being harmful and even carcinogenic to the human body and animals.

How? Studies maintain that CrO ₄ ^2- ions cross cell membranes by the action of sulfate-transporting proteins (both ions are in fact similar in size).

Reducing agents within cells reduce Cr (VI) to Cr (III), which accumulates by irreversibly coordinating to specific sites on macromolecules (such as DNA).

Once the cell is contaminated by an excess of chromium, it cannot leave due to the lack of mechanism that transports it back through the membranes.

Chromium uses

As colorant or pigments

Chromium has a wide range of applications, from colorant for different types of fabrics, to protector that embellishes metal parts in what is known as chrome plating, which can be made with pure metal, or with Cr (III) compounds or Cr (VI).

Chromic fluoride (CrF ₃), for example, is used as a colorant for woolen cloths; chromic sulfate (Cr ₂ (SO ₄) ₃), is used to color enamels, ceramics, paints, inks, varnishes, and also serves to chrome metals; and chromic oxide (Cr ₂ O ₃) also finds use where its attractive green color is required.

Therefore, any chromium mineral with intense colors can be destined to stain a structure, but after that the fact arises whether these compounds are dangerous or not for the environment or for the health of individuals.

In fact, its poisonous properties are used to preserve wood and other surfaces from insect attack.

In chrome or metallurgy

Small amounts of chromium are also added to the steel to strengthen it against oxidation and to improve its shine. This is due to the fact that it is capable of forming greyish carbides (Cr ₃ C ₂) that are very resistant when reacting with oxygen in the air.

Because chrome can be polished to glossy surfaces, chrome then features silver designs and colors as a cheaper alternative for these purposes.

Nutritional

Some debate whether chromium can be considered an essential element, that is, indispensable in the daily diet. It is present in some foods in very small concentrations, such as green leaves and tomatoes.

Likewise, there are protein supplements that regulate insulin activity and promote muscle growth, as is the case with chromium polynicotinate.

Where is it located?

Source: Pixabay

Chromium is found in a wide variety of minerals and gems such as rubies and emeralds. The main mineral from which chromium is extracted is chromite (MCr ₂ O ₄), where M can be any other metal with which chromium oxide is associated. These mines abound in South Africa, India, Turkey, Finland, Brazil and other countries.

Each source has one or more chromite variants. In this way, for each M (Fe, Mg, Mn, Zn, etc.) a different chromium mineral arises.

To extract the metal it is necessary to reduce the mineral, that is, to make the chromium metal center gain electrons by the action of a reducing agent. This is done with carbon or aluminum:

FeCr ₂ O ₄ + 4C => Fe + 2Cr + 4CO

Also, chromite (PbCrO ₄) is found.

Generally, in any mineral where the Cr ³⁺ ion can substitute for Al ³⁺, both with slightly similar ionic radii, it constitutes an impurity that results in another natural source of this amazing, but harmful, metal.

References

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4. NV Mandich. (nineteen ninety five). Chemistry of Chromium.. Taken from: citeseerx.ist.psu.edu
5. Chemistry LibreTexts. Chemistry of Chromium. Taken from: chem.libretexts.org
6. Saul 1. Shupack. (1991). The Chemistry of Chromium and Some Resulting Analytical Problems. Reviewed from: ncbi.nlm.nih.gov
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