Polyaluminum Chloride (PAC)

Polyaluminum chloride (polyaluminum chloride, abbreviated PAC; aluminum chlorohydrate), which is called polyaluminum. between AlCl3.

and Al(OH) 3. This combination of molecular formula [Al 2 (OH) n Cl 6-n · x H 2 O] m (m≤10, n=1~5), is a new type of flocculant polymeric mineral.

Table of Contents

1 Nature
2 Hazard information
3 Product use
4 Production, storage and transportation

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Nature

The pure product is yellow or gray flaky, granular or solid powder. It easily dissolves in water and the aqueous solution undergoes hydrolysis, which is accompanied by processes such as coagulation, absorption and precipitation.
Risk information

Non-flammable and irritating. Corrosive to tissues such as mucous membranes, eyes and skin.
It is dangerous for the environment.
Avoid contact with water, humid air, alkalis, alcohols, and combustible materials. Store in a cool, dry and ventilated place. Keep away from fire and heat sources.
Employees must use filter dust masks and protective glasses, acid and alkali resistant rubber clothes and acid and alkali resistant rubber gloves.
LD50: 316 mg/kg (oral in mice)

Use of the product

Polyaluminum chloride can be used as an antiperspirant in cosmetics, coagulant for water treatment, and also for industrial wastewater treatment.

production storage

• Polyaluminum chloride production methods mainly include aluminum ash hydrochloric acid method and boiling pyrolysis method
• In the meantime, the aluminum ash hydrochloric acid method is mainly adding aluminum ash to the washing water reactor and then adding hydrochloric acid for reaction and then sedimentation.
• To obtain the oxidation polymerization of liquid aluminum, finally obtained the solid polymerized alumina through dilution, filtration and drying.
• Polyaluminum chloride should be protected from exposure to rain and sun during transportation, and should be stored in a cool, ventilated and dry place, and the storage area should be clean.

Aluminum Chlorohydrate

Aluminum Chloride by Danny S

Aluminum chlorohydrate is a group of salts with the general formula Al n Cl (3n-m) (OH) m, which is used in deodorants.

Structure of polyaluminum chloride (PAC)

Aluminum chlorohydrate is one of the inorganic polymers whose structure is difficult to determine. However, recently scientific research groups, including Nazar and Laden

It depends on the Al 13 units. They concluded that the structure according to the arrangement of Kegen ions.

 Polyaluminum chloride (PAC) ready

Aluminum chlorohydrate is commercially prepared from the reaction of aluminum with hydrochloric acid.

Use

It is mainly used in deodorants.
Aluminum chlorohydrate is used in water purification as a floating substance, which is often called polyaluminum chloride. Due to the possibility of pH value to a specific value by changing the value of m and n in the structural formula is preferred.

Aluminum mono stearate

References LC Russell and LF Nazar. “Specification and Thermal Transformation in Alumina Salts: Structures of Polyhydroxy Aluminum Oxide Cluster [Al30O8(OH)56(H2O)26]18+ and its d-Keggin Part”. jam chemistry. Soc. (2000), 122 (15), 3777.
Carl Laden and Carl B. Felger. Antiperspirant and deodorant (cosmetic science and technology).

Aluminum hydroxy chloride

Medical Information
ATC-Code	M05 BX02
Specifications
Molar masses	It is not clear because of the mixture of ingredients
physical state	celebration [2]
Density	1.33-1.35 g·cm – 3 [3]
Dissolvability	Easily soluble in water: 500 g/L-1 (20°C) [4]
Safety instructions
Please note exemptions from labeling requirements for drugs, medical devices, cosmetics, food and animal feed. GHS labeling of hazardous substances [4] for 20-30% aqueous solution the danger H and P expressions H : 290-318 _ P .: 234-280-310-390-305 + 351 + 338 [ 5 ] _
Toxicological data	13 g kg-1 (TD Lo, mouse, female, pregnant, oral, continued 7-19 days) [6]
As far as possible and customary, SI units are used. Unless otherwise stated, the data given applies under standard conditions.

Aluminum hydroxychloride is a mixture of salts consisting of aluminum (Al), chlorine (Cl) and hydroxide (OH) with the composition Al n Cl (3n−m) (OH) m, for example Al 2 Cl(OH) 5 . They are usually produced and used as a mixture of inseparable single compounds, for example in polyaluminum chloride (PAC) cosmetics and body care products that reduce perspiration, as well as in wastewater treatment.

• 1 use
  • 1.1 flocculant
  • 1.2 Antiperspirant
• 2 Effects and risks
• 3. Toxicology
  • 3.1 Neurotoxicity
  • 3.2 Studies on the risk of breast cancer from aluminum hydroxychloride
• 4 See also
• 5. Individual proof

Use of polyaluminum chloride (PAC)

Strong, aluminum chloride hexahydrate is used in the textile and soapmaking industries for its astringent effect, where it is used to produce disinfectants or deodorants. Against mild inflammation in the throat area, solutions containing aluminum chloride or aluminum chlorate are provided for gargling. It is available in pharmacies without a prescription.

Aluminum hydroxychloride is classified as a high volume chemical. The mixture of these materials is used in the paper and textile industry and in water purification as a coagulant and sedimenting agent and in the cosmetic and health industries as an antiperspirant. [3]

Flocculant

Aluminum chloride containing water forms polymer structures called polyaluminum chloride (PAC). Considered an effective coagulant combination, it becomes a flocculant capable of converting a range of solutes into an undissolved state and removing many different types of suspended matter from aqueous solutions. Polymeric aluminum chlorides have been used since the 1970s in the treatment of service water, drinking water, wastewater, and swimming pool water as coagulants and sediments. In Europe, due to its technical and economic features, PAC replaced aluminum sulfate, which was common in drinking water treatment at that time. However, outside of Europe, distribution is low.

antiperspirant

Aluminum hydroxychloride can be used in certain concentrations by using it topically to prevent excessive sweating, and therefore it is also used as an active ingredient in many deodorants and antiperspirants. Aluminum chloride is corrosive, but according to experimental reports, polyaluminum chloride (PAC) may cause only minor skin irritation when the correct dosage is used. To reduce this problem, most antiperspirants contain glycerin or plant extracts.

The effect and dangers of polyaluminum chloride (PAC)

Aluminum chloride hexahydrate narrows pores by removing water and partially denatures proteins in skin cells and thus reduces sweating. Aluminum hydroxychloride can cause skin irritation, inflammation of the glands as a side effect, and granuloma. It is possible to develop eczema (“deodorant eczema”) and develop a permanent allergic reaction. As the effects of aluminum hydroxychloride in deodorant, long-term observations have shown mild signs of skin irritation. [7]

Toxicology of polyaluminum chloride (PAC)

Aluminum hydroxychloride was included by the European Union in 2014 in accordance with Regulation (EC) No. 1907/2006 (REACH) as part of the substance evaluation in the Community Roller Action Program (CoRAP). Here, the effects of this substance on human health and the environment are re-evaluated, and if necessary, follow-up measures are initiated. The reasons for using aluminum hydroxychloride were concerns about high tonnage (consolidation) and risks from possible assignment to the CMR material group. The re-evaluation started in 2015 and is carried out by France. More information was requested in order to reach a final assessment. [8]

Neurotoxicity of polyaluminum chloride (PAC)

Aluminum chloride damages the nervous system. [9] [10] [11] [12] At high doses, aluminum hydroxychloride disrupts the blood-brain barrier, can damage DNA, and has negative epigenetic effects. High doses of aluminum hydroxychloride have adverse effects on a number of species such as mammals, [14] rats, [15] rabbits, [9] and dogs. [16]

In February 2020, the Federal Institute for Risk Assessment (BfR) published a statement showing that human skin absorbs significantly less aluminum hydroxychloride than previously thought, especially through antiperspirants.

Studies on breast cancer risk from aluminum hydroxychloride

Aluminum salts such as aluminum chloride, aluminum zirconium tetrachlorohydrate complexes (“Aluminum Zirconium Tetrachlorohydrex Gly”) and aluminum hydroxy chloride in antiperspirant deodorants are suspected of causing breast cancer. On the one hand, there is an accumulation in the outer upper quadrant of the chest, which is close to the place where the deodorant is applied. There is more, however, epithelial tissue, which is a preferred site for cancer. In addition, increased aluminum concentration was found in polyaluminum chloride (PAC) samples of female breast cancer tissue. However, the association with the development of breast tumors was unclear and uptake into cells was unclear. [21] A 2008 meta-study that summarized previous research on the topic concluded that there was no scientific evidence to support this theory.

 Polyaluminum Chloride (PAC) in 2012

The Austrian Cancer Society asked Wolfram Parzefall (former University Professor of Toxicology at the Institute for Cancer Research at the Medical University of Vienna) to assess the carcinogenic (carcinogenic) risk of aluminum chloride (hexahydrate) as a component of deodorants. Since polyaluminum chloride (PAC) a previous publication (Sappino et al. 2012) [24] suggested a possible association with female breast cancer in vitro. This laboratory study with human breast cell cultures pointed out the cell destruction effect of aluminum chloride. The cells showed abnormal behavior comparable to the first phase of a tumor-like change. The aluminum chlorides used in this study were injected directly into the cell culture. The natural barrier of human skin was not considered.

Parzefall’s review notes that the American Cancer Society has published a more cautious assessment, citing aluminum compounds that alter estrogen receptors. These can be absorbed through the skin and lead to changes in estrogen receptors in breast cells. Because estrogen can cause the growth of both cancerous and non-cancerous breast cells, some scientists have suggested that the use of aluminum-based compounds in antiperspirants may be a risk factor for developing breast cancer. However, since no clear link to breast cancer has been established, researchers will continue to monitor aluminum hydroxychloride as a possible breast cancer risk factor. More studies are needed to draw clearer conclusions. In general, according to the Austrian Agency for Health and Food Safety (AGES) [26], it can be said that due to the different results, more research is needed to better understand the absorption of aluminum hydroxychloride. After skin application, the possible role of aluminum hydroxychloride in breast cell changes is determined. In terms of preventive health protection, such cosmetic products should not be placed in freshly shaved armpits.

In a reassessment from 2014

The BfR described the status of the study as contradictory. Antiperspirants containing aluminum help the absorption of aluminum in the human body. It is likely that a portion of the population has reached the tolerable weekly intake of 1 mg of aluminum per kilogram of body weight through food and other products containing aluminum. The BfR therefore recommends avoiding deodorants containing aluminum so as not to exceed the maximum tolerable limit. Despite a number of relevant studies, due to the state of conflicting data, a causal link between increased intake of aluminum hydroxychloride through antiperspirants and Alzheimer’s disease or breast cancer has not yet been scientifically proven. In 2019, the BfR confirmed that “it cannot be proven that aluminum is causally responsible for causing cancer”. The BfR also confirmed that, according to the current state of research, aluminum is neither genotoxic nor carcinogenic.

Recent studies have exaggerated the potential risks. For example, the European Union’s Scientific Committee on Consumer Safety is very clear that antiperspirants and cosmetics containing aluminum should be considered safe. The team, led by Hans Drexler, an environmental physician, also found in the study that aluminum-containing deodorants hardly absorb aluminum, at least over a two-week period. The BfR followed up on this assessment in 2020.

hydrogen bond

A hydrogen bond is a highly electrostatic force of attraction between an electronegative atom and a hydrogen atom to which another electronegative atom is covalently bonded. It results from the formation of a charge-dipole force with a hydrogen atom attached to a nitrogen, oxygen, or fluorine atom (hence the name “hydrogen bond”), not to be confused with covalent bonding to hydrogen atoms. The energy of a hydrogen bond (typically 5–30 kJ/mol) is significantly lower than that of weak covalent bonds (155 kJ/mol), and a typical covalent bond is only 20 times stronger than an intermolecular hydrogen bond. These bonds can be created between molecules (intermolecular) or between different parts of a molecule (intramolecular). 2 Hydrogen bonding is a very strong dipole-dipole electrostatic force that binds many molecules together because it gives great stability, but weaker than covalent bonding or ionic covalent bonding. ) Placed. This type of bond occurs both in inorganic molecules such as water and in organic molecules such as DNA.

Intermolecular hydrogen bonding is responsible for the high boiling point of water (100°C). This is because of the strong hydrogen bonding of chalcogen hydrides, unlike others. Intramolecular hydrogen bonding is partially responsible for secondary structure, tertiary structure, quaternary structure, and proteins and nucleic acids.

link

A hydrogen atom bonded to a relatively electronegative atom is a hydrogen bond donor atom. 3 This electronegative atom can be fluorine, oxygen or nitrogen. An electronegative atom such as fluorine, oxygen, or nitrogen is a hydrogen bond acceptor, regardless of whether it is covalently bonded to a hydrogen atom or not. An example of a hydrogen bond donor is ethanol, which has a hydrogen atom covalently bonded to oxygen. An example of a hydrogen bond acceptor with which the hydrogen diatom is not covalently bonded is the oxygen atom in ethyl ether.

Chloroform carbon can also participate in hydrogen bonding, when the carbon atom bonds to some electronegative atom, as in the case of CHCl 3 . The electronegative atom pulls the electron cloud around the hydrogen nucleus, and by decentralizing the cloud, it leaves the atom with a slight positive charge. Due to the small size of hydrogen compared to other atoms and molecules, the resulting charge, even if it is partial, shows a high charge density. A hydrogen bond is formed when this strong positive charge density attracts an electron pair from another heteroatom, which becomes a hydrogen bond acceptor.

Hydrogen bonding is usually described as an electrostatic dipole-dipole interaction.

However, it also has some of the characteristics of a covalent bond: it is directional, strong, creates interatomic distances less than the sum of the van der Waals radii, and usually contains a limited number of interaction partners that can be interpreted as a form of valence. . These covalent properties are important when acceptors bind hydrogen atoms from more electronegative donors.

The partially covalent nature of the hydrogen bond raises the question: “Which molecule does the hydrogen nucleus belong to?” And “Which one should be the “donor” and which one should be the “receiver”?” In general, this is easy to determine simply based on the interatomic distances of the XH … Y system: typically, the XH distance is ~1.1 Å, while the H … Y distance is ~1.6 to 2.0 Å. Liquids that show hydrogen bonding are called association liquids.

Hydrogen bonds can vary in strength, from very weak (1-2 kJ mol -1 >155 kJ mol -1, such as the HF2- ion) to very strong ( ). 4 Some typical values ​​are:

• F—H … F (155 kJ/mol)
• O—H … N (29 kJ/mol)
• O—H…O (21 kJ/mol)
• N—H … N (13 kJ/mol)
• N—H…O (8 kJ/mol)
• + OH 3 … :OH 2 (18 kJ/mol 5 ) (data obtained using molecular dynamics as mentioned in reference, and should be compared with 7.9 kJ/mol for raw water, also using dynamics A similar molecule has been obtained.

The length of hydrogen bonds depends on the strength of the bond. The bond strength itself depends on temperature, pressure, bond angle, and environment (typically characterized by the local dielectric constant). The typical hydrogen bond length in water is 1.97 Å (197 pm). The ideal bond angle depends on the nature of the hydrogen bond donor. Experimental results of hydrogen fluoride donor with different receptors show the following angles: 6

History

In his book The Nature of the Chemical Bond, Linus Pauling attributes the first mention of hydrogen bonding to TS Moore and TF Winmill in 1912 (J. Chem. Soc. 101, 1635). Moore and Winmill used hydrogen bonding to explain the fact that trimethylammonium hydroxide is a weaker base than tetramethylammonium hydroxide. The description of hydrogen bonding in its best known form, in water, was given a few years later, in 1920 by Latimer and Rudbush (JACS, 42, 1419).

Hydrogen bonding in water

The most common example of hydrogen bonding is water.

In an isolated water molecule, there are two hydrogen atoms and one oxygen atom. Two water molecules can form a hydrogen bond between them. In its simplest form, when there are only two molecules, it is called a water dimer and is often used as a model system. The more molecules there are, like in liquid water, the more bonds are possible because the oxygen in a water molecule has two lone pairs of electrons, each of which can form hydrogen bonds with carbon atoms. hydrogen from the other two atoms. Water Molecules This can be repeated, so that each water molecule is hydrogen bonded to up to four other water molecules, as shown in the figure (two through lone pairs, and two through hydrogen atoms. self).

The high boiling point of water is due to the large number of hydrogen bonds that each molecule has in relation to its low molecular mass, and due to the high strength of these hydrogen bonds. Water has very high boiling, melting and viscosity points, compared to other substances that are not connected by hydrogen bonds. The reason for these features is the difficulty of breaking these bonds. Water is unique in that its oxygen atoms have two lone pairs and two hydrogen atoms, meaning that the total number of bonds in a water molecule is four. For example, hydrogen fluoride—which has three lone pairs on the fluorine atom, but only one hydrogen atom—can only have two pairs in total. Ammonia has the opposite problem: three hydrogen atoms, but only one lone pair.

HF… HF… HF

The exact number of hydrogen bonds in which a molecule participates in liquid water

It fluctuates over time and depends on temperature. From simulations of TIP4P liquid water at 25 °C, it is estimated that each water molecule participates in an average of 3.59 hydrogen bonds. At 100°C, this number decreases to 3.24 due to increased molecular motion and consequent reduction in density, while at 0°C, the average number of hydrogen bonds increases to 3.69. 7 A recent study showed a much lower number of hydrogen bonds: 2357 at 25°C 8 These differences may be due to the use of a different method for defining and counting hydrogen bonds.

Where the bond strength is more equal, the atoms of two water molecules can be divided into two polyatomic ions of opposite charge, especially hydroxide (OH- and hydronium (H3O+)). (Hydronium ions are also known as hydroxonium ions.)

HO – H 3 O +

However, in pure water under normal conditions of pressure and temperature, this last formulation is rarely applicable. On average, every 5.5 × 108 molecules donate a proton to another water molecule, according to the water dissociation constant under such conditions. This is an important part of the uniqueness of water.

Bifurcated and hypercoordinated hydrogen bonds in water

It may happen that one hydrogen atom participates in two hydrogen bonds instead of one. This type of link is called “compressed”. It has been proposed that bifurcation hydrogen bonding is an essential step in water reorientation. 9

Hydrogen bond acceptors (which form branches with lone pairs of oxygen atoms (actually ending in “oxygen”)) are more abundant than donors 10

Hydrogen bonding in macromolecules and polymers

Hydrogen bonds determine the structure and properties of various macromolecular systems, both of natural and synthetic origin. Natural polymer macromolecules such as proteins (silk, spider silk, keratins, fibroins, etc.) or some structural polysaccharides such as cellulose or chitin are highly linked by hydrogen bonds. Likewise, the macromolecules of many synthetic polymers such as polyamides or polyurethanes are linked to varying degrees by hydrogen bonds. 11

Hydrogen bonds in DNA and proteins

Hydrogen bonding also plays an important role in determining the three-dimensional structures adopted by proteins and nucleic acids. In these macromolecules, the hydrogen bond between the parts of a molecule causes it to bend in a certain way, which helps to determine the physiological or biochemical role of the molecule. For example, the double helix structure of DNA is mainly due to the hydrogen bonding between base pairs that joins one complementary strand to another and allows for replication.

In proteins, hydrogen bonds are formed between skeletal oxygen atoms and amide hydrogen atoms. An alpha helix is ​​formed when the spacing of the amino acid residues involved in hydrogen bonding between positions i and i + 4 is regular. When the distance is smaller, between positions i and i + 3, a 3 10 helix is ​​formed. A beta sheet is formed when two strands are joined by hydrogen bonds involving alternating residues from each participating strand. Hydrogen bonds also participate in the formation of the tertiary structure of proteins through the interaction of R groups.

Symmetrical hydrogen bonding

A symmetric hydrogen bond is a special type of hydrogen bond in which the hydrogen nucleus is located exactly halfway between two atoms of the same element. The bond strength to each of these atoms is equal. This is an example of a three-center, two-electron bond. This type of bond is much stronger than “normal” hydrogen bonds. The effective bond order is 0.5, so its strength is comparable to covalent bonding. Many anhydrous acids such as hydrogen fluoride and formic acid have been seen in ice at high pressures as well as in the solid phase. It has also been seen in the bifluoride anion [FHF].

Symmetric hydrogen bonds have recently been observed spectroscopically in formic acid at high pressure (>GPa). Each hydrogen atom forms a partial covalent bond with two atoms instead of one. Symmetrical hydrogen bonding is assumed in ice at high pressures (Ice X). Low hydrogen bond barriers are formed when the distance between two heteroatoms is very small.

hydrogen bond

A hydrogen bond can be closely compared to a dihydrogen bond, which is also an intermolecular bonding interaction involving hydrogen atoms. These structures have been known for some time and have been well characterized by X-ray crystallography. However, understanding its relationship with conventional hydrogen bonding, ionic bonding, and covalent bonding is unclear. In general, hydrogen bonding is characterized by a proton acceptor, which is a lone pair of electrons in non-metal atoms (mainly nitrogen and oxygen). In some cases, these proton acceptors may be pi orbitals or multiple metal complexes. However, in a dihydrogen bond, a metal hydride acts as a proton acceptor. Forming a hydrogen-hydrogen interaction

Neutron diffraction has shown that the molecular geometry of these complexes is hydrogen bond-like, so that the bond length is suitable for metal/hydrogen complex donor systems.

Advanced hydrogen bonding theory

The nature of the link has recently been clarified. A widely published paper 12, from interpretations of anisotropy in the Compton profile of ordinary ice, proved that hydrogen bonding is partially covalent. Some nuclear magnetic resonance data on hydrogen bonding in proteins also indicate covalent bonding.

In general, hydrogen bonding can be viewed as a metric-dependent electrostatic scalar field between two or more intermolecular bonds. This is slightly different from intramolecular bound states, for example, covalent bonding or ionic bonding. However, hydrogen bonding is still a finite state phenomenon, as the interaction energy has a net negative sum. The early hydrogen bond theory proposed by Linus Pauling proposed that hydrogen bonds are partially covalent. This remained a controversial result until the late 1990s, when NMR techniques were used by F. Cordier et al. to transfer information between hydrogen-bonded nuclei, a feature that is only possible if the hydrogen bond contains some covalent properties.

Phenomena caused by hydrogen bonding

• Much higher boiling point of NH 3 , H 2 O and HF, compared to heavier analogs PH 3 , H 2 S and HCl.
• Viscosity of anhydrous phosphoric acid and glycerol.
• The formation of dimers in carboxylic acids and hexamers in hydrogen fluoride, which occurs even in the gas phase, and as a result, many deviations from the ideal gas law are created.
• The high solubility of many compounds such as ammonia in water is explained by hydrogen bonding with water molecules.
• Negative azeotropic mixture of HF and water.
• – NaOH precipitation is partly caused by the reaction of OH with moisture to form H 3 O 2 species – hydrogen bonds. A similar process occurs between NaNH 2 and NH 3 and between NaF and HF.
• The fact that ice is less dense than liquid water is due to its crystal structure, which is stabilized by hydrogen bonds.
• The presence of hydrogen bonds can cause anomalies in the natural sequence of aggregation states for certain mixtures of chemical compounds, with increasing or decreasing temperature. These compounds can be liquid up to a certain temperature, then solid even as the temperature increases, and finally liquid when the temperature rises above the “abnormal range”. 13
• Smart rubber uses hydrogen bonding as its only form of bonding, so it can “repair” itself when punctured, as new hydrogen bonds can appear between two surfaces of a polymer.

Hydrogen bond or hydrogen bond

Both hydrogen bond and hydrogen bond are correct but have very different meanings. In this case, the term that refers to the highly attractive electrostatic force between an electronegative atom and a hydrogen atom attached to another electronegative atom is hydrogen bonding. In both Spanish and English the term hydrogen bridge has become popular, but this name is incorrect and is not recommended, although it is still used in many textbooks.

According to IUPAC, it is written in English as hydrogen bond (hydrogen bond) and not hydrogen bond (bond with hydrogen bridge) or hydrogen bridge (hydrogen bridge). Therefore, a direct translation should be made to all other languages, including Spanish, and thus the correct term is hydrogen bonding.

The confusion between these two terms originated in English and spread to other languages. The term was later corrected in English, but in some languages, such as Spanish, the error was too deep. So much so that it is common to read recent textbooks that continue to misuse the term. In addition, the fact that a hydrogen bond is not a chemical bond, but an intermolecular bond, has favored the use of the word bridge, which avoids the confusion of thinking that a hydrogen bond is a chemical bond equivalent to a covalent bond. Uni – The word bridge is often used in high school textbooks so that students who are not familiar with chemistry terms do not confuse the concepts.

In fact, hydrogen bonding has nothing to do with hydrogen bonding.

Hydrogen bonds are unconventional covalent bonds that form to stabilize a particular compound. For example, in the case of boranes, the simplest boron (BH 3 ) is unstable, because the boron in this molecule can only have the expected six valence electrons and not the eight electrons expected in the second period elements, which are stabilized by having eight valence electrons. . As a result, boron trihydride tends to unite with another molecule to form diborane B 2 H 6 , where two hydrogen bonds are formed so that each boron has eight valence electrons and thus the noble gas configuration (in this case Neon) is available. Each hydrogen bond is formed by BHB, and unlike a normal covalent bond, where two electrons are used to connect two atoms (2c-2c), one pair of electrons is used to connect all three atoms. 3c-2e). Therefore, each boron receives the eight valence electrons it needs for stability. This is called “hydrogen bond” which has nothing to do with “hydrogen bond”.

The case of diborane is described, but there are many types of boranes with much more complex structures and many more types of bonds apart from hydrogen bonding. These are the various links that Burans can provide

• BH = 2c-2e terminal boron-hydrogen bond

• BHB = 3c-2e hydrogen bond
• BB = 2c-2e Enlace boro-boro
• BBB = 3c-2e bond with open boron bridge
• BBB = 3c-2e closed boron bond

The purpose of explaining the structure of boranes and their bonds is to make the reader understand that a hydrogen bond is of a very different nature than a hydrogen bond.

Covalent bond

A covalent bond occurs in two nonmetal atoms when they bond and share one or more electrons from the last level, an example of rolling a two and one on a dice is the first result of a logical sequence but not the total as twice the ones. (valence electrons) 1 (except for hydrogen, which is stabilized by having 2 electrons) so as to reach the octet rule. Electronegativity difference between atoms is not large enough for ionic bonding to occur. In order for a covalent bond to form, it is necessary that the electronegativity difference between the atoms is less than 1.7. 2

In this way, two atoms share one or more pairs of electrons in a new type of orbital called a molecular orbital. Covalent bonds are formed between atoms of a non-metallic element, between different non-metals and between a non-metal and hydrogen. 3 4

When non-metal atoms bond together in an ionic form, one is more electronegative than the other, so the electron cloud attracts the bond to its nucleus, creating an electric dipole. 5 This polarization allows the molecules of a compound to be attracted to each other by electrostatic forces of different strengths.

On the contrary, when the atoms of the same non-metallic element are covalently bonded, their electronegativity difference is zero and no dipole is created. Molecules have practically no attraction for each other.

In short, in an ionic bond, the transfer of electrons occurs from one atom to another, and in a covalent bond, the bonding electrons are shared by both atoms. In covalent bonding, two non-metal atoms share one or more electrons, that is, they join together through their electrons in the last orbital, which depends on the desired atomic number. Two atoms can share one, two, or three pairs of electrons, which leads to the formation of a single, double, or triple bond, respectively. In the Lewis structure, these bonds can be represented by a small line between the atoms.

History

The term “covalency” in relation to bonding was first coined in 1919 by Irving Langmuir in an article in the Journal of the American Chemical Society entitled “The Arrangement of Electrons in Atoms and Molecules.” ” used. In it, Langmuir wrote: “We use the term covalency to specify the number of pairs of electrons that a given atom shares with its neighbors.” 6

The idea of ​​covalent bonding can be traced back several years to Gilbert N. Lewis followed, describing in 1916 the exchange of electron pairs between atoms. 7 He introduced the Lewis symbol or electron dot symbol or Lewis dot structure, in which the valence electrons (those in the outer shell) are represented as dots around atomic symbols. The pairs of electrons between atoms represent covalent bonds. Multiple pairs represent multiple bonds such as double bonds and triple bonds. An alternative form of representation, not shown here, has electron bonding pairs shown as solid lines.

Lewis proposed that an atom forms enough covalent bonds to form a complete outer (or closed) electron shell. In the methane diagram shown here, the carbon atom has a valence of four and is therefore surrounded by eight electrons (the octet rule), four electrons from the carbon itself and four electrons from the hydrogens attached to it. Each hydrogen has a valence and is surrounded by two electrons (a binary rule), its own electron plus one from carbon. The number of electrons corresponds to the full shells in the quantum theory of the atom. The outer shell of the carbon atom is the n=2 shell, which holds eight electrons, while the outer (and only) shell of the hydrogen atom is the n=1 shell, which holds two or more elements.

While the idea of ​​shared pairs of electrons provides an effective qualitative picture of covalent bonding

Quantum mechanics is essential to understand the nature of these bonds and to predict the structures and properties of simple molecules. Walter Heitler and Fritz London provided the first successful explanation of a chemical bond using quantum mechanics, specifically molecular hydrogen, in 1927.8 Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is A good match between the atomic orbitals of the participating atoms.

These atomic orbitals have specific angular relationships to each other and thus the valence bond model can successfully predict the bond angles observed in simple molecules.

However, the theory of covalent bonding, or as the idea of ​​electron sharing based on the cubic atom, faced several conceptual problems, because this theory had competition with the ionic bond model. Despite this competition of these two theories, the covalent bond theory was accepted until 1920. M.Niaz and MARodríguez in their text Historia y filosofía de las ciencias: necesidad de su incorporationa en las textos Universityes de ciencias mention that Lewis recognizes that the cubic structure cannot represent the triple bond and suggests an alternative facet. Slow it down with four atoms. For years, Lewis hypothesized that if an atom’s electrons pair magnetically, it is easy to see how two unpaired electrons in different atoms can pair magnetically to form a nonpolar bond.

Types of covalent substances

There are two types of covalent substances:

Molecular covalent substances: covalent bonds form molecules that have the following characteristics: 9 10

• Low melting and boiling point.
• Under normal conditions of pressure and temperature (25 degrees Celsius) they can be solid, liquid or gas. ( About
• They are soft in the solid state.
• They are electrical current and heat insulators.
• Solubility: Polar molecules are soluble in polar solvents and non-polar molecules are soluble in non-polar solvents (such as similar solvents).
• For example: carbon dioxide, benzene, oxygen, nitrogen.

Lattice or Lattice Covalent Materials: In addition, covalent materials form crystalline networks of an unlimited number of atoms similar to ionic compounds that have these properties: 10

• High melting and boiling point.
• They are solid under normal conditions.
• They are very hard materials.
• They are insulators (except graphite).
• They are insoluble.
• Example: quartz, diamond.

Definition of covalent bond

Consider hydrogen atoms. When they are brought close together, the forces that attract each electron to the nucleus of the other atom become significant, until these attractive forces are neutralized by the repulsion that the electrons feel toward each other. At that point, the molecule presents the most stable configuration.

What has happened is that the orbitals of both electrons overlap, so that it is now impossible to tell which atom each electron belongs to.

According to chemists S. See and G. William Daub, in the hydrogen molecule, as in all covalent substances, four aspects must be considered:

First:

The properties of uncombined individual atoms are very different from the properties of molecules. For this reason, when the chemical formula of hydrogen is written, it should be written as H 2, because it is a diatomic molecule.

Second:

The two positive nuclei attract two electrons to produce a more stable molecule than the isolated atoms, resulting in a bond forming, resulting in a more stable molecule than the previous one. Because of the attraction that the nuclei exert on the two electrons, the repulsion between them is balanced, and therefore the electron is more likely to be found somewhere between the nuclei.

Third:

The distance between the nuclei should be such that the 1s orbitals overlap the most. In the case of the hydrogen molecule, the distance between the nuclei is approximately 0.74 Å. Otherwise, the distance between two covalently bonded atoms is called bond length.

Fourth:

52.0 kcal is required to “break” the covalent bonds in 1.0 g of hydrogen gas and form a hydrogen atom. 11

However, when the atoms are different, shared electrons are not equally attracted, so they tend to gravitate toward the most electronegative atom, that is, the atom with the greatest appetite for electrons. This phenomenon is called the polarity of polyaluminum chloride (PAC) (atoms with higher electronegativity acquire more negative polarity and pull the shared electrons towards their nucleus) and leads to the displacement of charges within the molecule.

It can be said that the most electronegative atom does not like to share its electrons with other atoms, and in the most extreme case, it wants the electron to be given to it unconditionally and thus an ionic bond is formed. Hence, it is said that polar covalent bonds have some ionic character.

When this difference is between 0 and 1.7, the covalent character is that which predominates in the case of the CH bond. However, according to chemist Raymond Chang, this electronegativity difference between atoms must be 2.0 or greater to be considered an ionic bond.Polyaluminum chloride (PAC) (Chang, 371).

Depending on the difference in electronegativity, covalent bonds can be classified as polar covalent and pure or nonpolar covalent. If the electronegativity difference is between 0.4 and 1.7, the covalent bond is polar, and if it is less than 0.4, the covalent bond is nonpolar.

When the electronegativity difference is zero (two equal atoms), the bond formed will be pure covalent. For an electronegativity difference of 1.9, the ionic character now reaches 35%, and for a difference of 3, it will be 49.5%.

Between oxygen or fluorine and elements of group 1 and 2, this difference will be maximum and its ionic characteristic will be as well.

Polar covalent bond

In a polar covalent bond, the electrons are shared unequally between the atoms and spend more time near one atom than the other. 12 Due to the unequal distribution of electrons between the atoms of different elements, slightly positive (δ+) and slightly negative (δ–) charges appear in different parts of the molecule.

In a water molecule, the bond that joins the oxygen to each hydrogen is a polar bond. Oxygen is a much more electronegative atom than hydrogen, so oxygen in water is partially negatively charged (high electron density), while hydrogens carry polyaluminum chloride (PAC) partially positive charges (high electron density). have a bottom).

In general, the relative electronegativity of two atoms in a bond, that is, their tendency to accumulate shared electrons, determines the polarity or nonpolarity of the bond. As long as one element is significantly more electronegative than the other, the bond between them will be polar. This means that one end will have a slightly positive charge and the other will have a slightly negative charge.

That is, it consists of the formation of bonds between the atoms of different elements, and the electronegativity difference must be greater than 0.4. In this bond, electrons are primarily attracted by the nucleus of the most electronegative atom, producing molecules whose electron cloud presents an area of ​​higher negative charge density and another area of ​​higher positive charge density (dipole).

non-polar covalent bond

Nonpolar covalent bonds are formed between two atoms of the same element or between atoms of different elements that share electrons more or less equally. 12 For example, molecular oxygen is nonpolar because the electrons are shared equally between the two oxygen atoms.

Another example of a non-polar covalent bond can be found in methane. Polyaluminum chloride (PAC) carbon has four electrons in its outer shell and needs four more electrons to become a stable octet. It achieves them by sharing electrons with four hydrogen atoms, each providing one electron. Similarly, hydrogen atoms each need an extra electron to fill their outermost shell, which they receive in the form of shared electrons from carbon. Although carbon and hydrogen do not have exactly the same electronegativity, they are quite similar, so carbon-hydrogen bonds are considered nonpolar.

Different types of covalent bonds

Simple link:

• A common electronic pair
• which is formed by an electron belonging to the last energy level of each atom
• and is indicated by a line. Example: HH, Cl-Cl

Double link:

• by two common electron pairs,
• That is, it is formed by two electrons belonging to the last energy level of each atom
• and is represented by two parallel lines. Example: OR=OR

• Triple bond: formed by three shared electronic pairs, i.e. polyaluminum chloride (PAC) by three electrons belonging to the last energy level of each atom and shown by three parallel lines. Example: N≡N

• Dative covalent bond or coordination: This is an electron pair shared by two atoms, but both electrons are formed by the same atom. It is usually indicated by an arrow (→).

An example of a chemical species that has a coordinate bond is ammonium ion (NH 4 1+). Ammonium ion consists of a proton and ammonia.

Compounds in which coordinate bonds are found are known as coordination compounds.

Coordinating compounds, also called complexes, usually bind to several surrounding anions known as ligands.

Ionic bond

Ionic bonding or 1 is the result of electrostatic attraction between ions of different sign according to the valence of the elements and the number of electrons that must be lost or gained to complete the layers, i.e. a layer. strongly electropositive and the other strongly electronegative. 2 This happens when in the bond, one of the atoms takes electrons from the other. The electrostatic attraction between the oppositely charged ions causes them to unite and form a simple chemical compound, where they do not melt. But one gives and the other takes. For an ionic bond to form, the difference (delta) of electronegativities must be greater than or equal to 1.7. (Pawling scale). 3

It should be noted that no bond is completely ionic, there will always be a contribution to the bond that can be attributed to the sharing of electrons in the same bond (covalency). 4 The ionic bond model is an exaggeration that is convenient because many thermodynamic data can be obtained with great accuracy if the atoms are treated as ions and there is no sharing of electrons.

Because the elements involved have large differences in electronegativity, this bond usually occurs between a metallic compound and a nonmetallic compound. 5 There is a complete electron transfer from one atom 6 to another, forming ions of different signs. The metal donates one or more electrons, creating positively charged ions or cations with a stable electronic configuration of negatively charged 7. These electrons then enter the nonmetal and form an ion or anion, which also has a stable electronic configuration of polyaluminum chloride (PAC). . They are stable because they both gain 8 electrons in their outermost shell (valence shell), according to the eight rule or Lewis structure, although this is not entirely true as we have several exceptions, such as hydrogen (H) which has an octet. has it.

Ions form lattices of crystalline compounds composed of oppositely charged N ions held together by electrostatic forces. This type of attraction determines the observed properties. If the electrostatic attraction is strong, crystalline solids with a high melting point and insoluble in water are formed. If the gravity is lower, such as NaCl, the melting point is also lower and they are generally soluble in water and insoluble in non-polar liquids such as benzene or carbon disulfide.

Features of polyaluminum chloride (PAC)

Some of the features of this type of link are:

• The bonds are very strong (depends strongly on the nature of the ions).
• It is solid at room temperature and has a crystalline or transparent structure in the cubic system. (There are ionic compounds that are liquid at room temperature called “ionic liquids” or “molten salts” with a gigantic field of application.)
• High melting point (between 300 °C and 1000 °C) and boiling point (if the bond has a high covalent character, these values ​​may decrease suddenly)
• They are bonds that result from interactions between Group I and II metals and Group VI and VII nonmetals.
• They and other solutions in water are soluble due to the electric dipole present in the water molecules. Polyaluminum chloride (PAC) is capable of solvating ions, thus compensating the crystal lattice energy. (Not all ionic compounds can be easily dissolved in water, either because of the low solvation energy of the ions or because of the covalent nature of the ionic compound).
• When placed in aqueous solution, they are excellent conductors of electricity, since then the ions are free. 9 (There is a wide variety of ionic compounds that are slightly or very slightly soluble in aqueous solution, also due to the covalent nature of this compound and does not allow water to easily separate the crystal lattice, thus resulting in a weak conductivity in solution)
• It has only simple links.
• In the solid state, they do not conduct electricity because the ions occupy very fixed positions in the lattice. If we use a block of salt as part of the circuit instead of the wire, the circuit will not work. To use, a light bulb does not work as part of the circuit. So, if from a bucket of water, but if we dissolve plenty of salt in said bucket, the circuit bulb lights up. This is due to the fact that the dissolved ions of the salt can go to the opposite pole (of their own sign) of the circuit battery and therefore, it works. 10

Classification of polyaluminum chloride (PAC)

There are two types of classifications:

A) Anion: It is an ion with a negative electric charge, which means that the atoms that make up polyaluminum chloride (PAC) have extra electrons. Anions are usually made of nonmetals, although certain anions are made of metals and nonmetals. The most common anions (number indicating charge):

b) Cation: It is an ion with a positive electrical charge. The most common ones are formed from metals, 12 but there are certain cations that are formed from non-metals.

Anion

The charge of the anion with the ion is more negative electric [ ], that is, the electrons. 2 Monoatomic anions have a negative oxidation state. Polyatomic anions are described as collections of bonded atoms with an overall negative electrical charge that change their individual oxidation states.

types of

There are three types of anions: monoatomic, polyatomic and acidic.

Polyatomic anions

They can be considered from a molecule that has gained an electron or from an acid that has lost a proton.

Traditional naming

They are named with the word ion or anion, followed by the nonmetallic name polyaluminum chloride (PAC), which ends in -ite for lower valencies or -ate for higher valences. Example:

Systematic nomenclature

They are named like acids but prefixed with the word ion or anion and omit the “hydrogen”. Example:

Acid anions

They come from a polyprotic acid that has lost some of its hydrogen atoms as protons. Polyprotic acids (or polybasic acids) are acids that have more than one ionizable hydrogen.

Traditional naming

They are named like the corresponding ion, but by adding the word acid and using multiplicative prefixes when there is more than one.

For diprotic acids (with two hydrogens in their formula) an old but deprecated nomenclature system is still maintained in commerce and industry. This includes naming the anion with the prefix bi-.

Systematic nomenclature

They are named as the corresponding ion but prefixed by hydrogen- with the corresponding multiplicative prefix.

For better understanding, we create a classification scheme, because it is not a rigid classification.

Classification scheme

Class (A)

They produce gases with dilute hydrochloric or sulfuric acid: carbonate, bicarbonate, sulfite, thiosulfate, sulfide, nitrite, hypochlorite, cyanide, and cyanate. Items (I) include fluoride, chloride, bromide, iodide, nitrate, chlorate, perchlorate, bromate and iodate, borate*, ferrocyanide, ferricyanide, thiocyanate, formate, acetate, oxalate, tartrate and citrate.

Class (B)

Precipitation reactions: sulfate, persulfate**, phosphate, phosphite, hypophosphite, arsenate, arsenite, silicate, fluorosilicate, salicylate, benzoate, and succinate. Oxidation and reduction reactions in solution: manganate, permanganate, chromate and dichromate.

common anions

Analytical march of the most common anions

Most common anions in the laboratory cannot be separated as well as cations. Most of the time they are detected directly, while others are separated into large groups that precipitate with cations and from these precipitates, anions are detected. However, it is much more difficult to analyze the anions present in the laboratory than the cations. In general, in the laboratory, the analytical implementation of anions is done by removing all existing cations by precipitation with sodium hydroxide or sodium carbonate. Then three primary tests are performed.

Salts are usually made of cations and anions (although the bond is never completely ionic, there is always a covalent contribution).

• acetate
• DNA is an anion.
• Many proteins are anionic at physiological pH.