HYDROGEN: HISTORY, STRUCTURE, PROPERTIES AND USES - CHEMISTRY - 2025

The hydrogen is a chemical element that is represented by the symbol H. The atom is the smallest of all and is the one that starts the periodic table, no matter where is positioned. It consists of a colorless gas made up of diatomic H ₂ molecules, not isolated H atoms; as with the noble gases He, Ne, Ar, among others.

Of all the elements, it is perhaps the most emblematic and outstanding, not only for its properties in terrestrial or drastic conditions, but for its immense abundance and variety of its compounds. Hydrogen is a gas, although inert in the absence of fire, flammable and dangerous; while water, H ₂ O, is the universal and life solvent.

Red cylinders used to store hydrogen. Source: Famartin

By itself, hydrogen does not show any visual peculiarities worthy of admiration, being simply a gas that is stored in cylinders or red bottles. However, it is its properties and ability to bond with all elements, which makes hydrogen special. And all this, despite the fact that it only has one valence electron.

If hydrogen were not stored in its respective cylinders, it would escape into space while much of it reacts on ascent. And although it has a very low concentration in the air we breathe, outside the Earth and in the rest of the Universe, it is the most abundant element, found in the stars and considered its unit of construction.

On Earth, on the other hand, it represents about 10% of its total mass. To visualize what this means, it must be considered that the surface of the planet is practically covered with oceans and that hydrogen is found in minerals, in crude oil and in any organic compound, in addition to being part of all living beings.

Like carbon, all biomolecules (carbohydrates, proteins, enzymes, DNA, etc.) have hydrogen atoms. Therefore, there are many sources to extract or produce it; however, few represent truly profitable production methods.

History

Identification and name

Although in 1671 Robert Boyle first witnessed a gas that was formed when iron filings reacted with acids, it was the British scientist Henry Cavendish, in 1766, who identified it as a new substance; the "flammable air".

Cavendish found that when this supposedly flammable air burned, water was generated. Based on his work and results, the French chemist Antoine Lavoisier gave this gas the name of hydrogen in 1783. Etymologically its meaning derives from the Greek words 'hydro' and 'genes': forming water.

Electrolysis and fuel

Soon after, in 1800, American scientists William Nicholson and Sir Anthony Carlisle discovered that water can decompose into hydrogen and oxygen; they had found the electrolysis of water. Later, in 1838, the Swiss chemist Christian Friedrich Schoenbein introduced the idea of ​​taking advantage of the combustion of hydrogen to generate electricity.

The popularity of hydrogen was such that even the writer Jules Verne referred to it as a fuel of the future in his book The Mysterious Island (1874).

Isolation

In 1899, the Scottish chemist James Dewar was the first to isolate hydrogen as a liquefied gas, being himself the one who was able to cool it enough to obtain it in its solid phase.

Two channels

From this point on, the history of hydrogen presents two channels. On the one hand, its development within the field of fuels and batteries; and on the other, the understanding of the structure of its atom and how it represented the element that opened the doors to quantum physics.

Structure and electronic configuration

Diatomic hydrogen molecule. Source: Benjah-bmm27

Hydrogen atoms are very small and have only one electron to form covalent bonds. When two of these atoms join, they give rise to a diatomic molecule, H ₂; this is molecular hydrogen gas (top image). Each white sphere corresponds to an individual H atom, and the global sphere to the molecular orbitals.

Thus, hydrogen actually consists of very small H ₂ molecules that interact through London scattering forces, since they lack a dipole moment because they are homonuclear. Therefore, they are very "restless" and spread rapidly in space as there are not strong enough intermolecular forces to slow them down.

The electron configuration of hydrogen is simply 1s ¹. This orbital, 1s, is the product of solving the famous Schrödinger equation for the hydrogen atom. In H _2, two 1s orbitals overlap to form two molecular orbitals: one bonding and the other anti-bonding, according to the molecular orbital theory (TOM).

These orbitals allow or explain the existence of the ions H ₂ ⁺ or H ₂ ^-; however, hydrogen chemistry is defined under normal conditions by H ₂ or H ⁺ or H ^- ions.

Oxidation numbers

From the electron configuration for hydrogen, 1s ¹, it is very easy to predict its possible oxidation numbers; bearing in mind, of course, that the higher-energy 2s orbital is not available for chemical bonds. Thus, in the basal state, hydrogen has an oxidation number of 0, H ⁰.

If it loses its only electron, the 1s orbital remains empty and the hydrogen cation or ion, H ⁺, is formed with great mobility in almost any liquid medium; especially the water. In this case, its oxidation number is +1.

And when the opposite happens, that is, gaining an electron, the orbital will now have two electrons and will become 1s ². Then the oxidation number becomes -1, and corresponds to the hydride anion, H ^-. It is worth noting that H ^- is isoelectronic to the noble gas helium, He; that is, both species have the same number of electrons.

In summary, the oxidation numbers for hydrogen are: +1, 0 and -1 and the molecule of H ₂ has as having two hydrogen atoms H ⁰.

Phases

The preferred phase of hydrogen, at least in terrestrial conditions, is the gaseous one, due to the reasons previously exposed. However, when temperatures decrease in the order of -200 ° C, or if the pressure increases hundreds of thousands of times than atmospheric, hydrogen can condense or crystallize into a liquid or solid phase, respectively.

Under these conditions, H ₂ molecules can be aligned in different ways to define structural patterns. The London scattering forces now become highly directional and therefore geometries or symmetries adopted by H ₂ pairs appear.

For example, two pairs H ₂, it is that equal to writing (H ₂) ₂ define a symmetric or asymmetric square. Meanwhile, three H ₂, or (H ₂) ₃ pairs define a hexagon, very similar to those of carbon in graphite crystals. In fact, this hexagonal phase is the main or most stable phase for solid hydrogen.

But what if the solid was made up not of molecules but of H atoms? Then we would deal with metallic hydrogen. These H atoms, recalling the white spheres, can define both a liquid phase and a metallic solid.

Properties

Physical appearance

Hydrogen is a colorless, odorless, and tasteless gas. Therefore, having a leak represents a risk of explosion.

Boiling point

-253 ° C.

Melting point

-259 ° C.

Flash point and stability

It explodes at virtually any temperature if there is a spark or heat source close to the gas, even sunlight can ignite hydrogen. However, as long as it is well stored it is a poorly reactive gas.

Density

0.082 g / L. It is 14 times lighter than air.

Solubility

1.62 mg / L at 21 ºC in water. It is, generally speaking, insoluble in most liquids.

Vapor pressure

1.24 · 10 ⁶ mmHg at 25 ° C. This value gives an idea of ​​how closed the hydrogen cylinders must be to prevent gas from escaping.

Autoignition temperature

560v ° C.

Electronegativity

2.20 on the Pauling scale.

Heat of combustion

-285.8 kJ / mol.

Heat of vaporization

0.90 kJ / mol.

Heat of fusion

0.117 kJ / mol.

Isotopes

The “normal” hydrogen atom is protium, ¹ H, which makes up about 99.985% of hydrogen. The other two isotopes for this element are deuterium, ² H, and tritium, ³ H. These differ in the number of neutrons; deuterium has one neutron, while tritium has two.

Spin isomers

There are two types of molecular hydrogen, H ₂: ortho and para. In the first, the two spins (of the proton) of the H atoms are oriented in the same direction (they are parallel); while in the second, the two spins are in opposite directions (they are antiparallel).

Hydrogen-para is the more stable of the two isomers; But as the temperature increases, the ortho: para ratio becomes 3: 1, which means that the hydrogen-ortho isomer predominates over the other. At very low temperatures (remotely close to absolute zero, 20K), pure hydrogen-para samples can be obtained.

Nomenclature

The nomenclature to refer to hydrogen is one of the simplest; although it is not the same way for its inorganic or organic compounds. H ₂ can be called by the following names in addition to 'hydrogen':

-Molecular hydrogen

-Dihydrogen

-Diatomic hydrogen molecule.

For the H ⁺ ion their names are proton or hydrogen ion; and if it is in an aqueous medium, H ₃ O ⁺, hydronium cation. While the H ^- ion is the hydride anion.

The hydrogen atom

The hydrogen atom represented by Bohr's planetary model. Source: Pixabay.

The hydrogen atom is the simplest of all and is normally represented as in the image above: a nucleus with a single proton (for ¹ H), surrounded by an electron that draws an orbit. All the atomic orbitals for the other elements of the periodic table have been constructed and estimated on this atom.

A more faithful representation to the current understanding of atoms would be that of a sphere whose periphery is defined by the electron and probabilistic cloud of the electron (its 1s orbital).

Where to find and production

A field of stars: inexhaustible source of hydrogen. Source: Pixabay.

Hydrogen is, though perhaps to a lesser degree compared to carbon, the chemical element that can be said without doubt to be everywhere; in the air, forming part of the water that fills the seas, oceans and our bodies, in crude oil and minerals, as well as in the organic compounds that are assembled to give rise to life.

Just skim any library of compounds to find hydrogen atoms in them.

The question is not so much how much but how it is present. For example, the molecule H ₂ is so volatile and reactive in the incidence of sunlight, which is very low in the atmosphere; therefore, it reacts to join other elements and thus gain stability.

While higher up in the cosmos, hydrogen is predominantly found as neutral atoms, H.

In fact, hydrogen is considered, in its metallic and condensed phase, as the building unit of stars. As there are immeasurable quantities of them and, due to its robustness and colossal dimensions, they make this element the most abundant in the entire universe. It is estimated that 75% of the known matter corresponds to hydrogen atoms.

natural

Collecting hydrogen atoms loose in space sounds impractical and extracting them from the peripheries of the Sun, or nebulae, unreachable. On Earth, where its conditions force this element to exist as H ₂, it can be produced through natural or geological processes.

For example, hydrogen has its own natural cycle in which certain bacteria, microbes, and algae can generate it through photochemical reactions. The scaling of natural processes and parallel to these includes the use of bioreactors, where bacteria feed on hydrocarbons to release the hydrogen contained in them.

Living things are also hydrogen producers, but to a lesser degree. If this were not the case, it would not be possible to explain how it constitutes one of the gaseous components of flatulence; which have been excessively proven to be flammable.

Finally, it should be mentioned that under anaerobic conditions (without oxygen), for example in underground layers, minerals can react slowly with water to produce hydrogen. Fayelita's reaction proves it:

3Fe ₂ SiO ₄ + 2 H ₂ O → 2 Fe ₃ O ₄ + 3 SiO ₂ + 3 H ₂

Industrial

Although biohydrogen is an alternative to generate this gas on an industrial scale, the most used methods practically consist of "removing" the hydrogen from the compounds that contain it, so that its atoms unite and form H ₂.

The least environmentally friendly methods of producing it consist of reacting coke (or charcoal) with superheated steam:

C (s) + H ₂ O (g) → CO (g) + H ₂ (g)

Likewise, natural gas has been used for this purpose:

CH ₄ (g) + H ₂ O (g) → CO (g) + 3H ₂ (g)

And because the amounts of coke or natural gas are vast, it is profitable to produce hydrogen by either of these two reactions.

Another method of obtaining hydrogen is to apply an electrical discharge to water to break it down into its elemental parts (electrolysis):

2 H ₂ O (l) → 2 H ₂ (g) + O ₂ (g)

At the laboratory

Molecular hydrogen can be prepared in small quantities in any laboratory. To do this, an active metal must be reacted with a strong acid, either in a beaker or in a test tube. The observable bubbling is a clear sign of hydrogen formation, represented by the following general equation:

M (s) + nH ⁺ (aq) → M ^{n +} (aq) + H ₂ (g)

Where n is the valence of the metal. For example, magnesium reacts with H ⁺ to produce H ₂:

Mg (s) + 2H ⁺ (aq) → Mg ²⁺ (aq) + H ₂ (g)

Reactions

Redox

The oxidation numbers by themselves offer a first glimpse of how hydrogen participates in chemical reactions. The H ₂ when reacting can remain unchanged, or split into the H ⁺ or H ^- ions depending on which species it binds with; if they are more or less electronegative than it.

H ₂ is not very reactive due to the strength of its covalent bond, HH; however, this is not an absolute impediment for it to react and form compounds with almost every element in the periodic table.

Its best known reaction is with that of oxygen gas to produce water vapors:

H ₂ (g) + O ₂ (g) → 2H ₂ O (g)

And such is its affinity for oxygen to form the stable water molecule, that it can even react with it as an O ^2- anion in certain metal oxides:

H ₂ (g) + CuO (s) → Cu (s) + H ₂ O (l)

Silver oxide also reacts or is "reduced" by the same reaction:

H ₂ (g) + AgO (s) → Ag (s) + H ₂ O (l)

These hydrogen reactions correspond to the redox type. That is, reduction-oxidation. Hydrogen oxidizes both in the presence of oxygen and of the metal oxides of metals less reactive than it; for example, copper, silver, tungsten, mercury, and gold.

Absorption

Some metals can absorb hydrogen gas to form metal hydrides, which are considered to be alloys. For example, transition metals such as palladium absorb significant amounts of H _2, being similar to metallic sponges.

The same happens with more complex metal alloys. In this way hydrogen can be stored by means other than its cylinders.

Addition

Organic molecules can also "absorb" hydrogen through different molecular mechanisms and / or interactions.

For metals, H ₂ molecules are surrounded by metal atoms within their crystals; while in organic molecules, the HH bond breaks to form other covalent bonds. In a more formalized sense: hydrogen is not absorbed, but is added to the structure.

The classic example is the addition of H ₂ to the double or triple bond of alkenes or alkynes, respectively:

C = C + H ₂ → HCCH

C≡C + H ₂ → HC = CH

These reactions are also called hydrogenation.

Hydride formation

Hydrogen reacts directly with elements to form a family of chemical compounds called hydrides. They are mainly of two types: saline and molecular.

Likewise, there are the metallic hydrides, which consist of the metallic alloys already mentioned when these metals absorb hydrogen gas; and the polymeric ones, with networks or chains of bonds EH, where E denotes the chemical element.

Saline

In saline hydrides, hydrogen participates in ionic bonding as the hydride anion, H ^-. For this to form, the element necessarily has to be less electronegative; otherwise, it would not give up its electrons to hydrogen.

Therefore, salt hydrides are only formed when hydrogen reacts with highly electropositive metals, such as alkali and alkaline earth metals.

For example, hydrogen reacts with metallic sodium to produce sodium hydride:

2Na (s) + H ₂ (g) → 2NaH (s)

Or with barium to produce barium hydride:

Ba (s) + H ₂ (g) → BaH ₂ (s)

Molecular

Molecular hydrides are even better known than ionic ones. They are also called hydrogen halides, HX, when hydrogen reacts with a halogen:

Cl ₂ (g) + H ₂ (g) → 2HCl (g)

Here hydrogen participates in the covalent bond as H ⁺; since, the differences between the electronegativities between both atoms is not very great.

Water itself can be considered as an oxygen hydride (or hydrogen oxide), the formation reaction of which has already been discussed. The reaction with sulfur is very similar to give hydrogen sulfide, a smelly gas:

S (s) + H ₂ (g) → H ₂ S (g)

But of all the molecular hydrides the most famous (and perhaps the most difficult to synthesize) is ammonia:

N ₂ (g) + 3H ₂ (g) → 2NH ₃ (g)

Applications

In the previous section, one of the main uses of hydrogen was already addressed: as a raw material for the development of synthesis, inorganic or organic. Controlling this gas usually has no other purpose than to make it react to create compounds other than those from which it was extracted.

Raw material

- It is one of the reagents for the synthesis of ammonia, which in turn has endless industrial applications, starting with the production of fertilizers, even as a material to nitrogenate drugs.

- It is intended to react with carbon monoxide and thus massively produce methanol, a reagent that is highly important in biofuels.

Reducing agent

- It is a reducing agent for certain metal oxides, which is why it is used in metallurgical reduction (already explained in the case of copper and other metals).

- Reduce fats or oils to produce margarine.

Oil industry

In the oil industry, hydrogen is used to "hydrotreat" crude oil in refining processes.

For example, it seeks to fragment large and heavy molecules into small molecules with greater demand in the market (hydrocracking); release the metals trapped in the petroporphyrin cages (hydrodemetallization); remove sulfur atoms as H ₂ S (hydrodesulfurization); or reduce double bonds to create paraffin-rich mixtures.

Fuel

Hydrogen itself is an excellent fuel for rockets or spacecraft, since small amounts of it, when reacting with oxygen, release enormous amounts of heat or energy.

On a smaller scale, this reaction is used to design hydrogen cells or batteries. However, these cells face the difficulties of not being able to store this gas properly; and the challenge of becoming completely independent from burning fossil fuels.

On the positive side, hydrogen used as fuel releases only water; instead of gases that represent means of pollution for the atmosphere and ecosystems.

References

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