1.
Organosulfate
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Organosulfates are a class of organic compounds sharing a common functional group commonly with the structure R-O-SO3−. The SO4 core is a group and the R group is any organic residue. All organosulfates are formally derived from alcohols and sulfuric acid. Many sulfate esters are used in detergents, and some are useful reagents, alkyl sulfates consist of a hydrophobic hydrocarbon chain, a polar sulfate group and either a cation or amine to neutralize the sulfate group. Examples include, sodium sulfate and related potassium and ammonium salts. Alkyl sulfates are commonly used as a surfactant in liquid soaps, detergents to clean wool, surface cleaners, as active ingredients in laundry detergents, shampoos. They can also be found in household products such as toothpastes, antacids, cosmetics. Generally they are found in consumer products at concentrations ranging from 3-20%, in 2003 approximately 118,000 t/a of alkyl sulfates were used in the U. S. A common example is sodium lauryl sulfate, with the formula CH311OSO3Na, also common in consumer products are the sulfate esters of ethoxylated fatty alcohols such as those derived from lauryl alcohol. An example is sodium laureth sulfate, an ingredient in some cosmetics, alkylsulfate can be produced from alcohols, which in turn are obtained by hydrogenation of animal or vegetable oils and fats or using the Ziegler process or through oxo synthesis. If produced from oleochemical feedstock or the Ziegler process, the chain of the alcohol will be linear. If derived using the oxo process, a low level of branching will appear usually with a methyl or ethyl group at the C-2 position, containing even and odd amounts of alkyl chains. These alcohols react with acid, ClSO3H + ROH → ROSO3H + HCl Some organosulfates can be prepared by the Elbs persulfate oxidation of phenols. A less common family of organosulfates have the formula R-O-SO2-O-R and they are prepared from sulfuric acid and the alcohol. The main examples are diethyl sulfate and dimethyl sulfate, colourless liquids that are used as reagents in organic synthesis and these compounds are potentially dangerous alkylating agents. The reduction of sulfate in nature involves the formation of one or sometimes two sulfate esters, adenosine 5-phosphosulfate and 3-phosphoadenosine-5-phosphosulfate, sulfate is an inert anion, so nature activates it by the formation of these ester derivatives, which are susceptible to reduction to sulfite. Many organisms utilize these reactions for metabolic purposes or for the biosynthesis of sulfur compounds required for life, because they are widely used in commercial products, the safety aspects of organosulfates are heavily investigated. Alkyl sulfates if ingested are well-absorbed and are metabolized into a C3, C4 or C5 sulfate, the highest irritant of the alkyl sulfates is sodium laurylsulfate, with the threshold before irritation at a concentration of 20%
2.
Jmol
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Jmol is computer software for molecular modelling chemical structures in 3-dimensions. Jmol returns a 3D representation of a molecule that may be used as a teaching tool and it is written in the programming language Java, so it can run on the operating systems Windows, macOS, Linux, and Unix, if Java is installed. It is free and open-source software released under a GNU Lesser General Public License version 2.0, a standalone application and a software development kit exist that can be integrated into other Java applications, such as Bioclipse and Taverna. A popular feature is an applet that can be integrated into web pages to display molecules in a variety of ways, for example, molecules can be displayed as ball-and-stick models, space-filling models, ribbon diagrams, etc. Jmol supports a range of chemical file formats, including Protein Data Bank, Crystallographic Information File, MDL Molfile. There is also a JavaScript-only version, JSmol, that can be used on computers with no Java, the Jmol applet, among other abilities, offers an alternative to the Chime plug-in, which is no longer under active development. While Jmol has many features that Chime lacks, it does not claim to reproduce all Chime functions, most notably, Chime requires plug-in installation and Internet Explorer 6.0 or Firefox 2.0 on Microsoft Windows, or Netscape Communicator 4.8 on Mac OS9. Jmol requires Java installation and operates on a variety of platforms. For example, Jmol is fully functional in Mozilla Firefox, Internet Explorer, Opera, Google Chrome, fast and Scriptable Molecular Graphics in Web Browsers without Java3D
3.
ChemSpider
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ChemSpider is a database of chemicals. ChemSpider is owned by the Royal Society of Chemistry, the database contains information on more than 50 million molecules from over 500 data sources including, Each chemical is given a unique identifier, which forms part of a corresponding URL. This is an approach to develop an online chemistry database. The search can be used to widen or restrict already found results, structure searching on mobile devices can be done using free apps for iOS and for the Android. The ChemSpider database has been used in combination with text mining as the basis of document markup. The result is a system between chemistry documents and information look-up via ChemSpider into over 150 data sources. ChemSpider was acquired by the Royal Society of Chemistry in May,2009, prior to the acquisition by RSC, ChemSpider was controlled by a private corporation, ChemZoo Inc. The system was first launched in March 2007 in a release form. ChemSpider has expanded the generic support of a database to include support of the Wikipedia chemical structure collection via their WiChempedia implementation. A number of services are available online. SyntheticPages is an interactive database of synthetic chemistry procedures operated by the Royal Society of Chemistry. Users submit synthetic procedures which they have conducted themselves for publication on the site and these procedures may be original works, but they are more often based on literature reactions. Citations to the published procedure are made where appropriate. They are checked by an editor before posting. The pages do not undergo formal peer-review like a journal article. The comments are moderated by scientific editors. The intention is to collect practical experience of how to conduct useful chemical synthesis in the lab, while experimental methods published in an ordinary academic journal are listed formally and concisely, the procedures in ChemSpider SyntheticPages are given with more practical detail. Comments by submitters are included as well, other publications with comparable amounts of detail include Organic Syntheses and Inorganic Syntheses
4.
European Chemicals Agency
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ECHA is the driving force among regulatory authorities in implementing the EUs chemicals legislation. ECHA helps companies to comply with the legislation, advances the safe use of chemicals, provides information on chemicals and it is located in Helsinki, Finland. The Agency, headed by Executive Director Geert Dancet, started working on 1 June 2007, the REACH Regulation requires companies to provide information on the hazards, risks and safe use of chemical substances that they manufacture or import. Companies register this information with ECHA and it is freely available on their website. So far, thousands of the most hazardous and the most commonly used substances have been registered, the information is technical but gives detail on the impact of each chemical on people and the environment. This also gives European consumers the right to ask whether the goods they buy contain dangerous substances. The Classification, Labelling and Packaging Regulation introduces a globally harmonised system for classifying and labelling chemicals into the EU. This worldwide system makes it easier for workers and consumers to know the effects of chemicals, companies need to notify ECHA of the classification and labelling of their chemicals. So far, ECHA has received over 5 million notifications for more than 100000 substances, the information is freely available on their website. Consumers can check chemicals in the products they use, Biocidal products include, for example, insect repellents and disinfectants used in hospitals. The Biocidal Products Regulation ensures that there is information about these products so that consumers can use them safely. ECHA is responsible for implementing the regulation, the law on Prior Informed Consent sets guidelines for the export and import of hazardous chemicals. Through this mechanism, countries due to hazardous chemicals are informed in advance and have the possibility of rejecting their import. Substances that may have effects on human health and the environment are identified as Substances of Very High Concern 1. These are mainly substances which cause cancer, mutation or are toxic to reproduction as well as substances which persist in the body or the environment, other substances considered as SVHCs include, for example, endocrine disrupting chemicals. Companies manufacturing or importing articles containing these substances in a concentration above 0 and they are required to inform users about the presence of the substance and therefore how to use it safely. Consumers have the right to ask the retailer whether these substances are present in the products they buy, once a substance has been officially identified in the EU as being of very high concern, it will be added to a list. This list is available on ECHA’s website and shows consumers and industry which chemicals are identified as SVHCs, Substances placed on the Candidate List can then move to another list
5.
PubChem
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PubChem is a database of chemical molecules and their activities against biological assays. The system is maintained by the National Center for Biotechnology Information, a component of the National Library of Medicine, PubChem can be accessed for free through a web user interface. Millions of compound structures and descriptive datasets can be downloaded via FTP. PubChem contains substance descriptions and small molecules with fewer than 1000 atoms and 1000 bonds, more than 80 database vendors contribute to the growing PubChem database. PubChem consists of three dynamically growing primary databases, as of 28 January 2016, Compounds,82.6 million entries, contains pure and characterized chemical compounds. Substances,198 million entries, contains also mixtures, extracts, complexes, bioAssay, bioactivity results from 1.1 million high-throughput screening programs with several million values. PubChem contains its own online molecule editor with SMILES/SMARTS and InChI support that allows the import and export of all common chemical file formats to search for structures and fragments. In the text search form the database fields can be searched by adding the name in square brackets to the search term. A numeric range is represented by two separated by a colon. The search terms and field names are case-insensitive, parentheses and the logical operators AND, OR, and NOT can be used. AND is assumed if no operator is used, example,0,5000,50,10 -5,5 PubChem was released in 2004. The American Chemical Society has raised concerns about the publicly supported PubChem database and they have a strong interest in the issue since the Chemical Abstracts Service generates a large percentage of the societys revenue. To advocate their position against the PubChem database, ACS has actively lobbied the US Congress, soon after PubChems creation, the American Chemical Society lobbied U. S. Congress to restrict the operation of PubChem, which they asserted competes with their Chemical Abstracts Service
6.
International Chemical Identifier
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Initially developed by IUPAC and NIST from 2000 to 2005, the format and algorithms are non-proprietary. The continuing development of the standard has supported since 2010 by the not-for-profit InChI Trust. The current version is 1.04 and was released in September 2011, prior to 1.04, the software was freely available under the open source LGPL license, but it now uses a custom license called IUPAC-InChI Trust License. Not all layers have to be provided, for instance, the layer can be omitted if that type of information is not relevant to the particular application. InChIs can thus be seen as akin to a general and extremely formalized version of IUPAC names and they can express more information than the simpler SMILES notation and differ in that every structure has a unique InChI string, which is important in database applications. Information about the 3-dimensional coordinates of atoms is not represented in InChI, the InChI algorithm converts input structural information into a unique InChI identifier in a three-step process, normalization, canonicalization, and serialization. The InChIKey, sometimes referred to as a hashed InChI, is a fixed length condensed digital representation of the InChI that is not human-understandable. The InChIKey specification was released in September 2007 in order to facilitate web searches for chemical compounds and it should be noted that, unlike the InChI, the InChIKey is not unique, though collisions can be calculated to be very rare, they happen. In January 2009 the final 1.02 version of the InChI software was released and this provided a means to generate so called standard InChI, which does not allow for user selectable options in dealing with the stereochemistry and tautomeric layers of the InChI string. The standard InChIKey is then the hashed version of the standard InChI string, the standard InChI will simplify comparison of InChI strings and keys generated by different groups, and subsequently accessed via diverse sources such as databases and web resources. Every InChI starts with the string InChI= followed by the version number and this is followed by the letter S for standard InChIs. The remaining information is structured as a sequence of layers and sub-layers, the layers and sub-layers are separated by the delimiter / and start with a characteristic prefix letter. The six layers with important sublayers are, Main layer Chemical formula and this is the only sublayer that must occur in every InChI. The atoms in the formula are numbered in sequence, this sublayer describes which atoms are connected by bonds to which other ones. Describes how many hydrogen atoms are connected to each of the other atoms, the condensed,27 character standard InChIKey is a hashed version of the full standard InChI, designed to allow for easy web searches of chemical compounds. Most chemical structures on the Web up to 2007 have been represented as GIF files, the full InChI turned out to be too lengthy for easy searching, and therefore the InChIKey was developed. With all databases currently having below 50 million structures, such duplication appears unlikely at present, a recent study more extensively studies the collision rate finding that the experimental collision rate is in agreement with the theoretical expectations. Example, Morphine has the structure shown on the right, as the InChI cannot be reconstructed from the InChIKey, an InChIKey always needs to be linked to the original InChI to get back to the original structure
7.
Simplified molecular-input line-entry system
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The simplified molecular-input line-entry system is a specification in form of a line notation for describing the structure of chemical species using short ASCII strings. SMILES strings can be imported by most molecule editors for conversion back into two-dimensional drawings or three-dimensional models of the molecules, the original SMILES specification was initiated in the 1980s. It has since modified and extended. In 2007, a standard called OpenSMILES was developed in the open-source chemistry community. Other linear notations include the Wiswesser Line Notation, ROSDAL and SLN, the original SMILES specification was initiated by David Weininger at the USEPA Mid-Continent Ecology Division Laboratory in Duluth in the 1980s. The Environmental Protection Agency funded the project to develop SMILES. It has since modified and extended by others, most notably by Daylight Chemical Information Systems. In 2007, a standard called OpenSMILES was developed by the Blue Obelisk open-source chemistry community. Other linear notations include the Wiswesser Line Notation, ROSDAL and SLN, in July 2006, the IUPAC introduced the InChI as a standard for formula representation. SMILES is generally considered to have the advantage of being slightly more human-readable than InChI, the term SMILES refers to a line notation for encoding molecular structures and specific instances should strictly be called SMILES strings. However, the term SMILES is also used to refer to both a single SMILES string and a number of SMILES strings, the exact meaning is usually apparent from the context. The terms canonical and isomeric can lead to confusion when applied to SMILES. The terms describe different attributes of SMILES strings and are not mutually exclusive, typically, a number of equally valid SMILES strings can be written for a molecule. For example, CCO, OCC and CC all specify the structure of ethanol, algorithms have been developed to generate the same SMILES string for a given molecule, of the many possible strings, these algorithms choose only one of them. This SMILES is unique for each structure, although dependent on the algorithm used to generate it. These algorithms first convert the SMILES to a representation of the molecular structure. A common application of canonical SMILES is indexing and ensuring uniqueness of molecules in a database, there is currently no systematic comparison across commercial software to test if such flaws exist in those packages. SMILES notation allows the specification of configuration at tetrahedral centers, and these are structural features that cannot be specified by connectivity alone and SMILES which encode this information are termed isomeric SMILES
8.
Chemical formula
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These are limited to a single typographic line of symbols, which may include subscripts and superscripts. A chemical formula is not a name, and it contains no words. Although a chemical formula may imply certain simple chemical structures, it is not the same as a full chemical structural formula. Chemical formulas can fully specify the structure of only the simplest of molecules and chemical substances, the simplest types of chemical formulas are called empirical formulas, which use letters and numbers indicating the numerical proportions of atoms of each type. Molecular formulas indicate the numbers of each type of atom in a molecule. For example, the formula for glucose is CH2O, while its molecular formula is C6H12O6. This is possible if the relevant bonding is easy to show in one dimension, an example is the condensed molecular/chemical formula for ethanol, which is CH3-CH2-OH or CH3CH2OH. For reasons of structural complexity, there is no condensed chemical formula that specifies glucose, chemical formulas may be used in chemical equations to describe chemical reactions and other chemical transformations, such as the dissolving of ionic compounds into solution. A chemical formula identifies each constituent element by its chemical symbol, in empirical formulas, these proportions begin with a key element and then assign numbers of atoms of the other elements in the compound, as ratios to the key element. For molecular compounds, these numbers can all be expressed as whole numbers. For example, the formula of ethanol may be written C2H6O because the molecules of ethanol all contain two carbon atoms, six hydrogen atoms, and one oxygen atom. Some types of compounds, however, cannot be written with entirely whole-number empirical formulas. An example is boron carbide, whose formula of CBn is a variable non-whole number ratio with n ranging from over 4 to more than 6.5. When the chemical compound of the consists of simple molecules. These types of formulas are known as molecular formulas and condensed formulas. A molecular formula enumerates the number of atoms to reflect those in the molecule, so that the formula for glucose is C6H12O6 rather than the glucose empirical formula. However, except for very simple substances, molecular chemical formulas lack needed structural information, for simple molecules, a condensed formula is a type of chemical formula that may fully imply a correct structural formula. For example, ethanol may be represented by the chemical formula CH3CH2OH
9.
Conjugate acid
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A conjugate acid, within the Brønsted–Lowry acid–base theory, is a species formed by the reception of a proton by a base—in other words, it is a base with a hydrogen ion added to it. On the other hand, a base is merely what is left after an acid has donated a proton in a chemical reaction. Hence, a base is a species formed by the removal of a proton from an acid. A proton is a particle with a unit positive electrical charge, it is represented by the symbol H+ because it constitutes the nucleus of a hydrogen atom, that is. A cation can be an acid, and an anion can be a conjugate base, depending on which substance is involved. In an acid-base reaction, an acid plus a base reacts to form a conjugate base plus a conjugate acid, refer to the following figure, We say that the water molecule is the conjugate acid of the hydroxide ion after the latter received the hydrogen proton donated by ammonium. On the other hand, ammonia is the base for the acid ammonium after ammonium has donated a hydrogen ion towards the production of the water molecule. The strength of an acid is directly proportional to its dissociation constant. If a conjugate acid is strong, its dissociation will have an equilibrium constant. The strength of a base can be seen as the tendency of the species to pull hydrogen protons towards itself. If a conjugate base is classified as strong, it will hold on to the hydrogen proton when in solution, if a chemical species is classified as a weak acid, its conjugate base will be strong in nature. This can be observed in ammonias reaction with water, the reaction proceeds until most of the ammonia has been transformed to ammonium. This shift to the right in the equilibrium of the reaction means that ammonium does not dissociate easily in water. On the other hand, if a species is classified as a strong acid, an example of this case would be the dissociation of Hydrochloric acid HCl in water. Since HCl is an acid, its conjugate base will be a weak conjugate base. Therefore, in system, most H+ will be in the form of a Hydronium ion H 3O+ instead of attached to a Cl anion. To summarize, the stronger the acid or base, the weaker the conjugate, the acid and conjugate base as well as the base and conjugate acid are known as conjugate pairs. When finding a conjugate acid or base, it is important to look at the reactants of the chemical equation, to identify the conjugate acid, look for the pair of compounds that are related
10.
Sulfate
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The sulfate or sulphate ion is a polyatomic anion with the empirical formula SO2−4. Sulfate is the spelling recommended by IUPAC, but sulphate is used in British English, salts, acid derivatives, and peroxides of sulfate are widely used in industry. Sulfates occur widely in everyday life, sulfates are salts of sulfuric acid and many are prepared from that acid. The sulfate anion consists of a sulfur atom surrounded by four equivalent oxygen atoms in a tetrahedral arrangement. The symmetry is the same as that of methane, the sulfur atom is in the +6 oxidation state while the four oxygen atoms are each in the state. The sulfate ion carries a charge of −2 and it is the conjugate base of the bisulfate ion, HSO−4. Organic sulfate esters, such as sulfate, are covalent compounds. The tetrahedral molecular geometry of the ion is as predicted by VSEPR theory. Later, Linus Pauling used valence bond theory to propose that the most significant resonance canonicals had two pi bonds involving d orbitals and his reasoning was that the charge on sulfur was thus reduced, in accordance with his principle of electroneutrality. The S−O bond length of 149 pm is shorter than the bond lengths in sulfuric acid of 157 pm for S−OH, the double bonding was taken by Pauling to account for the shortness of the S−O bond. Paulings use of d orbitals provoked a debate on the importance of π bonding. The outcome was a consensus that d orbitals play a role. A widely accepted description involving pπ – dπ bonding was initially proposed by D. W. J. Cruickshank, in this model, fully occupied p orbitals on oxygen overlap with empty sulfur d orbitals. However, in description, despite there being some π character to the S−O bonds. For sulfuric acid, computational analysis confirms a positive charge on sulfur. Therefore, the representation with four single bonds is the optimal Lewis structure rather than the one with two double bonds, in this model, the structure obeys the octet rule and the charge distribution is in agreement with the electronegativity of the atoms. However, the representation of Pauling for sulfate and other main group compounds with oxygen is still a common way of representing the bonding in many textbooks. On the other hand, in the structure with an ionic bond, exceptions include calcium sulfate, strontium sulfate, lead sulfate, and barium sulfate, which are poorly soluble
11.
American and British English spelling differences
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Many of the differences between American and British English date back to a time when spelling standards had not yet developed. For instance, some spellings seen as American today were once used in Britain. But English-language spelling reform has rarely been adopted otherwise, and thus modern English orthography varies somewhat between countries and is far from phonemic in any country, in the early 18th century, English spelling was inconsistent. These differences became noticeable after the publishing of influential dictionaries, todays British English spellings mostly follow Johnsons A Dictionary of the English Language, while many American English spellings follow Websters An American Dictionary of the English Language. Webster was a proponent of English spelling reform for reasons both philological and nationalistic, in A Companion to the American Revolution, John Algeo notes, it is often assumed that characteristically American spellings were invented by Noah Webster. He was very influential in popularizing certain spellings in America, rather he chose already existing options such as center, color and check for the simplicity, analogy or etymology. William Shakespeares first folios, for example, used spellings like center and color as much as centre, Webster did attempt to introduce some reformed spellings, as did the Simplified Spelling Board in the early 20th century, but most were not adopted. In Britain, the influence of those who preferred the Norman spellings of words proved to be decisive, later spelling adjustments in the United Kingdom had little effect on todays American spellings and vice versa. For the most part, the systems of most Commonwealth countries. Australian spelling has also strayed slightly from British spelling, with some American spellings incorporated as standard, New Zealand spelling is almost identical to British spelling, except in the word fiord. There is also an increasing use of macrons in words that originated in Māori, most words ending in an unstressed -our in British English end in -or in American English. Wherever the vowel is unreduced in pronunciation, e. g. contour, velour, paramour and troubadour the spelling is the same everywhere, most words of this kind came from Latin, where the ending was spelled -or. They were first adopted into English from early Old French, after the Norman conquest of England, the ending became -our to match the Old French spelling. The -our ending was not only used in new English borrowings, however, -or was still sometimes found, and the first three folios of Shakespeares plays used both spellings before they were standardised to -our in the Fourth Folio of 1685. After the Renaissance, new borrowings from Latin were taken up with their original -or ending, Websters 1828 dictionary had only -or and is given much of the credit for the adoption of this form in the United States. Johnson, unlike Webster, was not an advocate of spelling reform and he preferred French over Latin spellings because, as he put it, the French generally supplied us. In Jeffersons original draft it is spelled honour, Honor and honour were equally frequent in Britain until the 17th century, honor still is, in the UK, the usual spelling as a persons name and appears in Honor Oak, a district of London. In derivatives and inflected forms of the words, British usage depends on the nature of the suffix used
12.
Polyatomic ion
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The prefix poly- means many, in Greek, but even ions of two atoms are commonly referred to as polyatomic. In older literature, an ion is also referred to as a radical. In contemporary usage, the term refers to free radicals that are species with an unpaired electron. An example of an ion is the hydroxide ion, consisting of one oxygen atom and one hydrogen atom. An ammonium ion is made up of one atom and four hydrogen atoms, it has a charge of +1. Polyatomic ions are often useful in the context of chemistry or in the formation of salts. A polyatomic ion can often be considered as the conjugate acid/base of a neutral molecule, for example, the conjugate base of sulfuric acid is the polyatomic hydrogen sulfate anion. The removal of hydrogen ion yields the sulfate anion. There are two rules that can be used for learning the nomenclature of polyatomic anions. First, when the prefix bi is added to a name, a hydrogen is added to the formula and its charge is increased by 1. An alternative to the bi- prefix is to use the hydrogen in its place. Note that many of the common polyatomic anions are conjugate bases of acids derived from the oxides of non-metallic elements, for example, the sulfate anion, SO42−, is derived from H 2SO4, which can be regarded as SO3 + H 2O. The second rule looks at the number of oxygens in an ion, consider the chlorine oxoanion family, First, think of the -ate ion as being the base name, in which case the addition of a per- prefix adds an oxygen. Changing the -ate suffix to -ite will reduce the oxygens by one, in all situations, the charge is not affected. The naming pattern follows within many different oxyanion series based on a root for that particular series. The -ite has one less oxygen than the -ate, but different -ate anions might have different numbers of oxygen atoms and these rules will not work with all polyatomic anions, but they do work with the most common ones. The following tables give examples of commonly encountered polyatomic ions, only a few representatives are given, as the number of polyatomic ions encountered in practice is very large. Monatomic ions Salt Mass spectrometry Hydrogen peroxide Molecule Hydrogen molecular ion List of polyatomic ions A Beginners Guide To Polyatomic Ions, tables of Common Polyatomic Ions, including PDB files
13.
Ion
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An ion is an atom or a molecule in which the total number of electrons is not equal to the total number of protons, giving the atom or molecule a net positive or negative electrical charge. Ions can be created, by chemical or physical means. In chemical terms, if an atom loses one or more electrons. If an atom gains electrons, it has a net charge and is known as an anion. Ions consisting of only a single atom are atomic or monatomic ions, because of their electric charges, cations and anions attract each other and readily form ionic compounds, such as salts. In the case of ionization of a medium, such as a gas, which are known as ion pairs are created by ion impact, and each pair consists of a free electron. The word ion comes from the Greek word ἰόν, ion, going and this term was introduced by English physicist and chemist Michael Faraday in 1834 for the then-unknown species that goes from one electrode to the other through an aqueous medium. Faraday also introduced the words anion for a charged ion. In Faradays nomenclature, cations were named because they were attracted to the cathode in a galvanic device, arrhenius explanation was that in forming a solution, the salt dissociates into Faradays ions. Arrhenius proposed that ions formed even in the absence of an electric current, ions in their gas-like state are highly reactive, and do not occur in large amounts on Earth, except in flames, lightning, electrical sparks, and other plasmas. These gas-like ions rapidly interact with ions of charge to give neutral molecules or ionic salts. These stabilized species are commonly found in the environment at low temperatures. A common example is the present in seawater, which are derived from the dissolved salts. Electrons, due to their mass and thus larger space-filling properties as matter waves, determine the size of atoms. Thus, anions are larger than the parent molecule or atom, as the excess electron repel each other, as such, in general, cations are smaller than the corresponding parent atom or molecule due to the smaller size of its electron cloud. One particular cation contains no electrons, and thus consists of a single proton - very much smaller than the parent hydrogen atom. Since the electric charge on a proton is equal in magnitude to the charge on an electron, an anion, from the Greek word ἄνω, meaning up, is an ion with more electrons than protons, giving it a net negative charge. A cation, from the Greek word κατά, meaning down, is an ion with fewer electrons than protons, there are additional names used for ions with multiple charges
14.
International Union of Pure and Applied Chemistry
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The International Union of Pure and Applied Chemistry /ˈaɪjuːpæk/ or /ˈjuːpæk/ is an international federation of National Adhering Organizations that represents chemists in individual countries. It is a member of the International Council for Science, IUPAC is registered in Zürich, Switzerland, and the administrative office, known as the IUPAC Secretariat, is in Research Triangle Park, North Carolina, United States. This administrative office is headed by IUPACs executive director, currently Lynn Soby, IUPAC was established in 1919 as the successor of the International Congress of Applied Chemistry for the advancement of chemistry. Its members, the National Adhering Organizations, can be national chemistry societies, national academies of sciences, there are fifty-four National Adhering Organizations and three Associate National Adhering Organizations. IUPACs Inter-divisional Committee on Nomenclature and Symbols is the world authority in developing standards for the naming of the chemical elements. Since its creation, IUPAC has been run by different committees with different responsibilities. These committees run different projects which include standardizing nomenclature, finding ways to bring chemistry to the world, IUPAC is best known for its works standardizing nomenclature in chemistry and other fields of science, but IUPAC has publications in many fields including chemistry, biology and physics. IUPAC is also known for standardizing the atomic weights of the elements through one of its oldest standing committees, the need for an international standard for chemistry was first addressed in 1860 by a committee headed by German scientist Friedrich August Kekulé von Stradonitz. This committee was the first international conference to create an international naming system for organic compounds, the ideas that were formulated in that conference evolved into the official IUPAC nomenclature of organic chemistry. IUPAC stands as a legacy of this meeting, making it one of the most important historical international collaborations of chemistry societies, since this time, IUPAC has been the official organization held with the responsibility of updating and maintaining official organic nomenclature. IUPAC as such was established in 1919, one notable country excluded from this early IUPAC is Germany. Germanys exclusion was a result of prejudice towards Germans by the Allied powers after World War I, Germany was finally admitted into IUPAC during 1929. However, Nazi Germany was removed from IUPAC during World War II, during World War II, IUPAC was affiliated with the Allied powers, but had little involvement during the war effort itself. After the war, East and West Germany were eventually readmitted to IUPAC, since World War II, IUPAC has been focused on standardizing nomenclature and methods in science without interruption. In 2016, IUPAC denounced the use of chlorine as a chemical weapon, the letter stated, Our organizations deplore the use of chlorine in this manner. According to the CWC, the use, stockpiling, distribution, IUPAC is governed by several committees that all have different responsibilities. Each committee is made up of members of different National Adhering Organizations from different countries, the steering committee hierarchy for IUPAC is as follows, All committees have an allotted budget to which they must adhere. Any committee may start a project, if a projects spending becomes too much for a committee to continue funding, it must take the issue to the Project Committee
15.
Salt (chemistry)
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In chemistry, a salt is an ionic compound that results from the neutralization reaction of an acid and a base. Salts are composed of related numbers of cations and anions so that the product is electrically neutral and these component ions can be inorganic, such as chloride, or organic, such as acetate, and can be monatomic, such as fluoride, or polyatomic, such as sulfate. There are several varieties of salts, salts that hydrolyze to produce hydroxide ions when dissolved in water are alkali salts, whilst those that hydrolyze to produce hydronium ions in water are acidic salts. Neutral salts are those salts that are neither acidic nor basic, zwitterions contain an anionic centre and a cationic centre in the same molecule, but are not considered to be salts. Examples of zwitterions include amino acids, many metabolites, peptides, usually, non-dissolved salts at standard conditions for temperature and pressure are solid, but there are exceptions. Molten salts and solutions containing dissolved salts are called electrolytes, as they are able to conduct electricity. As observed in the cytoplasm of cells, in blood, urine, plant saps and mineral waters, therefore, their salt content is given for the respective ions. Salts can appear to be clear and transparent, opaque, and even metallic, in many cases, the apparent opacity or transparency are only related to the difference in size of the individual monocrystals. Since light reflects from the boundaries, larger crystals tend to be transparent. The color of the salt is due to the electronic structure in the d-orbitals of transition elements or in the conjugated organic dye framework. Different salts can elicit all five basic tastes, e. g. salty, sweet, sour, bitter, and umami or savory. Salts of strong acids and strong bases are non-volatile and odorless and that slow, partial decomposition is usually accelerated by the presence of water, since hydrolysis is the other half of the reversible reaction equation of formation of weak salts. Many ionic compounds can be dissolved in water or other similar solvents, the exact combination of ions involved makes each compound have a unique solubility in any solvent. The solubility is dependent on how well each ion interacts with the solvent, for example, all salts of sodium, potassium and ammonium are soluble in water, as are all nitrates and many sulfates – barium sulfate, calcium sulfate and lead sulfate are examples of exceptions. However, ions that bind tightly to each other and form highly stable lattices are less soluble, for example, most carbonate salts are not soluble in water, such as lead carbonate and barium carbonate. Some soluble carbonate salts are, sodium carbonate, potassium carbonate, solid salts do not conduct electricity. Moreover, solutions of salts also conduct electricity, the name of a salt starts with the name of the cation followed by the name of the anion. Salts are often referred to only by the name of the cation or by the name of the anion. g
16.
Sulfuric acid
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Sulfuric acid is a highly corrosive strong mineral acid with the molecular formula H2SO4 and molecular weight 98.079 g/mol. It is a pungent-ethereal, colorless to slightly yellow liquid that is soluble in water at all concentrations. Sometimes, it is dyed dark brown during production to alert people to its hazards, the historical name of this acid is oil of vitriol. Sulfuric acid is an acid and shows different properties depending upon its concentration. Its corrosiveness on other materials, like metals, living tissues or even stones, can be ascribed to its strong acidic nature and, if concentrated. It is also hygroscopic, readily absorbing water vapour from the air, Sulfuric acid at a high concentration can cause very serious damage upon contact, since not only does it cause chemical burns via hydrolysis, but also secondary thermal burns through dehydration. It can lead to permanent blindness if splashed onto eyes and irreversible damage if swallowed, Sulfuric acid has a wide range of applications including in domestic acidic drain cleaners, as an electrolyte in lead-acid batteries and in various cleaning agents. It is also a central substance in the chemical industry, principal uses include mineral processing, fertilizer manufacturing, oil refining, wastewater processing, and chemical synthesis. It is widely produced with different methods, such as process, wet sulfuric acid process, lead chamber process. The study of vitriol, a category of glassy minerals from which the acid can be derived, sumerians had a list of types of vitriol that they classified according to the substances color. Some of the earliest discussions on the origin and properties of vitriol is in the works of the Greek physician Dioscorides, galen also discussed its medical use. Ibn Sina focused on its uses and different varieties of vitriol. Sulfuric acid was called oil of vitriol by medieval European alchemists because it was prepared by roasting green vitriol in an iron retort, there are references to it in the works of Vincent of Beauvais and in the Compositum de Compositis ascribed to Saint Albertus Magnus. A passage from Pseudo-Geber´s Summa Perfectionis was long considered to be the first recipe for sulfuric acid, in the seventeenth century, the German-Dutch chemist Johann Glauber prepared sulfuric acid by burning sulfur together with saltpeter, in the presence of steam. As saltpeter decomposes, it oxidizes the sulfur to SO3, in 1736, Joshua Ward, a London pharmacist, used this method to begin the first large-scale production of sulfuric acid. This process allowed the effective industrialization of sulfuric acid production, after several refinements, this method, called the lead chamber process or chamber process, remained the standard for sulfuric acid production for almost two centuries. Sulfuric acid created by John Roebucks process approached a 65% concentration, later refinements to the lead chamber process by French chemist Joseph Louis Gay-Lussac and British chemist John Glover improved concentration to 78%. However, the manufacture of dyes and other chemical processes require a more concentrated product
17.
Sulfur
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Sulfur or sulphur is a chemical element with symbol S and atomic number 16. It is abundant, multivalent, and nonmetallic, under normal conditions, sulfur atoms form cyclic octatomic molecules with chemical formula S8. Elemental sulfur is a yellow crystalline solid at room temperature. Chemically, sulfur reacts with all elements except for gold, platinum, iridium, tellurium, though sometimes found in pure, native form, sulfur usually occurs as sulfide and sulfate minerals. Being abundant in native form, sulfur was known in ancient times, being mentioned for its uses in ancient India, ancient Greece, China, in the Bible, sulfur is called brimstone. Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas, the greatest commercial use of the element is the production of sulfuric acid for sulfate and phosphate fertilizers, and other chemical processes. The element sulfur is used in matches, insecticides, and fungicides, many sulfur compounds are odoriferous, and the smells of odorized natural gas, skunk scent, grapefruit, and garlic are due to organosulfur compounds. Hydrogen sulfide gives the characteristic odor to rotting eggs and other biological processes, sulfur is an essential element for all life, but almost always in the form of organosulfur compounds or metal sulfides. Three amino acids and two vitamins are organosulfur compounds, many cofactors also contain sulfur including glutathione and thioredoxin and iron–sulfur proteins. Disulfides, S–S bonds, confer mechanical strength and insolubility of the keratin, found in outer skin, hair. Sulfur is one of the chemical elements needed for biochemical functioning and is an elemental macronutrient for all organisms. Sulfur is derived from the Latin word sulpur, which was Hellenized to sulphur, the spelling sulfur appears toward the end of the Classical period. In 12th-century Anglo-French, it was sulfre, in the 14th century the Latin ph was restored, for sulphre, the parallel f~ph spellings continued in Britain until the 19th century, when the word was standardized as sulphur. Sulfur was the form chosen in the United States, whereas Canada uses both, the IUPAC adopted the spelling sulfur in 1990, as did the Nomenclature Committee of the Royal Society of Chemistry in 1992, restoring the spelling sulfur to Britain. Sulfur forms polyatomic molecules with different chemical formulas, the best-known allotrope being octasulfur, cyclo-S8. The point group of cyclo-S8 is D4d and its dipole moment is 0 D. Octasulfur is a soft, bright-yellow solid that is odorless and it melts at 115.21 °C, boils at 444.6 °C and sublimes easily. At 95.2 °C, below its melting temperature, cyclo-octasulfur changes from α-octasulfur to the β-polymorph, the structure of the S8 ring is virtually unchanged by this phase change, which affects the intermolecular interactions. At higher temperatures, the viscosity decreases as depolymerization occurs, molten sulfur assumes a dark red color above 200 °C
18.
Atom
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An atom is the smallest constituent unit of ordinary matter that has the properties of a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms, Atoms are very small, typical sizes are around 100 picometers. Atoms are small enough that attempting to predict their behavior using classical physics - as if they were billiard balls, through the development of physics, atomic models have incorporated quantum principles to better explain and predict the behavior. Every atom is composed of a nucleus and one or more bound to the nucleus. The nucleus is made of one or more protons and typically a number of neutrons. Protons and neutrons are called nucleons, more than 99. 94% of an atoms mass is in the nucleus. The protons have an electric charge, the electrons have a negative electric charge. If the number of protons and electrons are equal, that atom is electrically neutral, if an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively, and it is called an ion. The electrons of an atom are attracted to the protons in a nucleus by this electromagnetic force. The number of protons in the nucleus defines to what chemical element the atom belongs, for example, the number of neutrons defines the isotope of the element. The number of influences the magnetic properties of an atom. Atoms can attach to one or more other atoms by chemical bonds to form compounds such as molecules. The ability of atoms to associate and dissociate is responsible for most of the changes observed in nature. The idea that matter is made up of units is a very old idea, appearing in many ancient cultures such as Greece. The word atom was coined by ancient Greek philosophers, however, these ideas were founded in philosophical and theological reasoning rather than evidence and experimentation. As a result, their views on what look like. They also could not convince everybody, so atomism was but one of a number of competing theories on the nature of matter. It was not until the 19th century that the idea was embraced and refined by scientists, in the early 1800s, John Dalton used the concept of atoms to explain why elements always react in ratios of small whole numbers
19.
Oxygen
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Oxygen is a chemical element with symbol O and atomic number 8. It is a member of the group on the periodic table and is a highly reactive nonmetal. By mass, oxygen is the third-most abundant element in the universe, after hydrogen, at standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. This is an important part of the atmosphere and diatomic oxygen gas constitutes 20. 8% of the Earths atmosphere, additionally, as oxides the element makes up almost half of the Earths crust. Most of the mass of living organisms is oxygen as a component of water, conversely, oxygen is continuously replenished by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide. Oxygen is too reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone, strongly absorbs ultraviolet UVB radiation, but ozone is a pollutant near the surface where it is a by-product of smog. At low earth orbit altitudes, sufficient atomic oxygen is present to cause corrosion of spacecraft, the name oxygen was coined in 1777 by Antoine Lavoisier, whose experiments with oxygen helped to discredit the then-popular phlogiston theory of combustion and corrosion. One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle, Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries later Leonardo da Vinci built on Philos work by observing that a portion of air is consumed during combustion and respiration, Oxygen was discovered by the Polish alchemist Sendivogius, who considered it the philosophers stone. In the late 17th century, Robert Boyle proved that air is necessary for combustion, English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. From this he surmised that nitroaereus is consumed in both respiration and combustion, Mayow observed that antimony increased in weight when heated, and inferred that the nitroaereus must have combined with it. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract De respiratione. Robert Hooke, Ole Borch, Mikhail Lomonosov, and Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element. This may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, which was then the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, one part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. The fact that a substance like wood gains overall weight in burning was hidden by the buoyancy of the combustion products
20.
Tetrahedron
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In geometry, a tetrahedron, also known as a triangular pyramid, is a polyhedron composed of four triangular faces, six straight edges, and four vertex corners. The tetrahedron is the simplest of all the ordinary convex polyhedra, the tetrahedron is the three-dimensional case of the more general concept of a Euclidean simplex. The tetrahedron is one kind of pyramid, which is a polyhedron with a polygon base. In the case of a tetrahedron the base is a triangle, like all convex polyhedra, a tetrahedron can be folded from a single sheet of paper. For any tetrahedron there exists a sphere on which all four vertices lie, a regular tetrahedron is one in which all four faces are equilateral triangles. It is one of the five regular Platonic solids, which have known since antiquity. In a regular tetrahedron, not only are all its faces the same size and shape, regular tetrahedra alone do not tessellate, but if alternated with regular octahedra they form the alternated cubic honeycomb, which is a tessellation. The regular tetrahedron is self-dual, which means that its dual is another regular tetrahedron, the compound figure comprising two such dual tetrahedra form a stellated octahedron or stella octangula. This form has Coxeter diagram and Schläfli symbol h, the tetrahedron in this case has edge length 2√2. Inverting these coordinates generates the dual tetrahedron, and the together form the stellated octahedron. In other words, if C is the centroid of the base and this follows from the fact that the medians of a triangle intersect at its centroid, and this point divides each of them in two segments, one of which is twice as long as the other. The vertices of a cube can be grouped into two groups of four, each forming a regular tetrahedron, the symmetries of a regular tetrahedron correspond to half of those of a cube, those that map the tetrahedra to themselves, and not to each other. The tetrahedron is the only Platonic solid that is not mapped to itself by point inversion, the regular tetrahedron has 24 isometries, forming the symmetry group Td, isomorphic to the symmetric group, S4. The first corresponds to the A2 Coxeter plane, the two skew perpendicular opposite edges of a regular tetrahedron define a set of parallel planes. When one of these intersects the tetrahedron the resulting cross section is a rectangle. When the intersecting plane is one of the edges the rectangle is long. When halfway between the two edges the intersection is a square, the aspect ratio of the rectangle reverses as you pass this halfway point. For the midpoint square intersection the resulting boundary line traverses every face of the tetrahedron similarly, if the tetrahedron is bisected on this plane, both halves become wedges
21.
Oxidation state
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The oxidation state, often called the oxidation number, is an indicator of the degree of oxidation of an atom in a chemical compound. This is never exactly true for real bonds, the term oxidation was first used by Lavoisier to signify reaction of a substance with oxygen. Much later, it was realized that the substance, upon being oxidized, loses electrons, oxidation states are typically represented by integers. In some cases, the oxidation state of an element is a fraction. It is predicted that even a +10 oxidation state may be achieveable by platinum in the platinum tetroxide dication, the increase in oxidation state of an atom, through a chemical reaction, is known as an oxidation, a decrease in oxidation state is known as a reduction. Such reactions involve the transfer of electrons, a net gain in electrons being a reduction. For pure elements, the state is zero. There are various methods for determining oxidation states, in inorganic nomenclature, the oxidation state is determined and expressed as an oxidation number, and is represented by a Roman numeral placed after the element name. In coordination chemistry, oxidation number is defined differently from oxidation state, a “Comprehensive definition of oxidation state ” has been published with free access. This article is a distillation of an IUPAC technical report Toward a comprehensive definition of oxidation state. Exceptions to this are that hydrogen has a state of −1 in hydrides of active metals, e. g. LiH. There are two different methods for determining the state of elements in chemical compounds. First a method based on the rules in the IUPAC definition, second, a method based on the relative electronegativity of the elements in the compound, where the more electronegative element is assumed to take the negative charge. Any pure element—even if it forms diatomic molecules like chlorine —has an oxidation state of zero, examples of this are Cu or O2. For monatomic ions, the state is the same as the charge of the ion. For example, the anion has an oxidation state of −2. The sum of oxidation states for all atoms in a molecule or polyatomic ion is equal to the charge of the molecule or ion, thus, the oxidation state of one element can be calculated from the oxidation states of the other elements. An application of this rule is that the sum of the states of all atoms in a neutral molecule must be zero
22.
Dimethyl sulfate
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Dimethyl sulfate is a chemical compound with formula 2SO2. As the diester of methanol and sulfuric acid, its formula is written as 2SO4 or even Me2SO4. Me2SO4 is mainly used as an agent in organic synthesis. Under standard conditions, Me2SO4 is an oily liquid with a slight onion-like odour. Like all strong alkylating agents, Me2SO4 is extremely toxic and its use as a laboratory reagent has been superseded to some extent by methyl triflate, CF3SO3CH3, the methyl ester of trifluoromethanesulfonic acid. Dimethyl sulfate was first discovered in the early 19th century in an impure form, P Claesson later extensively studied its preparation. It was used in warfare in WW I. A common process is the reaction of dimethyl ether with sulfur trioxide. 2O + SO3 → 2SO4 Dimethyl sulfate is best known as a reagent for the methylation of phenols, amines, typically, one methyl group is transferred more quickly than the second. Methyl transfer is typically assumed to occur via an SN2 reaction, compared to other methylating agents, dimethyl sulfate is preferred by the industry because of its low cost and high reactivity. Most commonly, Me2SO4 is employed to methylate phenols and this process can be used to determine base sequencing, cleavage on the DNA chain, and other applications. Dimethyl sulfate also methylates adenine in single-stranded portions of DNA, upon re-annealing, these methyl groups interfere with adenine-guanine base-pairing. Nuclease S1 can then be used to cut the DNA in single-stranded regions and this is an important technique for analyzing protein-DNA interactions. Although dimethyl sulfate is effective and affordable, its toxicity has encouraged the use of other methylating reagents. Methyl iodide is a reagent used for O-methylation, like dimethyl sulfate, Dimethyl carbonate, which is less reactive, has far lower toxicity compared to both dimethyl sulfate and methyl iodide and can be used instead of dimethyl sulfate for N-methylation. High pressure can be used to accelerate methylation by dimethyl carbonate, in general, the toxicity of methylating agents is correlated with their efficiency as methyl transfer reagents. Dimethyl sulfate is carcinogenic and mutagenic, highly poisonous, corrosive, environmentally hazardous and it is considered a potential chemical weapon. Dimethyl sulfate is absorbed through the skin, mucous membranes, and gastrointestinal tract, an ocular reaction is also common
23.
Ester
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In chemistry, esters are chemical compounds derived from an acid in which at least one –OH group is replaced by an –O–alkyl group. Usually, esters are derived from an acid and an alcohol. Glycerides, which are fatty acid esters of glycerol, are important esters in biology, being one of the classes of lipids. Esters with low weight are commonly used as fragrances and found in essential oils. Phosphoesters form the backbone of DNA molecules, nitrate esters, such as nitroglycerin, are known for their explosive properties, while polyesters are important plastics, with monomers linked by ester moieties. The word ester was coined in 1848 by German chemist Leopold Gmelin, probably as a contraction of the German Essigäther, ester names are derived from the parent alcohol and the parent acid, where the latter may be organic or inorganic. Esters derived from more complex carboxylic acids are, on the hand, more frequently named using the systematic IUPAC name. For example, the ester hexyl octanoate, also known under the trivial name hexyl caprylate, has the formula CH36CO25CH3, the chemical formulas of organic esters usually take the form RCO2R′, where R and R′ are the hydrocarbon parts of the carboxylic acid and the alcohol, respectively. For example, butyl acetate, derived from butanol and acetic acid would be written CH3CO2C4H9, alternative presentations are common including BuOAc and CH3COOC4H9. Cyclic esters are called lactones, regardless of whether they are derived from an organic or an inorganic acid, one example of a lactone is γ-valerolactone. An uncommon class of organic esters are the orthoesters, which have the formula RC3, triethylorthoformate is derived, in terms of its name from orthoformic acid and ethanol. Esters can also be derived from an acid and an alcohol. For example, triphenyl phosphate is the derived from phosphoric acid. Organic carbonates are derived from acid, for example, ethylene carbonate is derived from carbonic acid. So far an alcohol and inorganic acid are linked via oxygen atoms, in corollary, boron features borinic esters, boronic esters, and borates. As oxygen is a group 16 chemical element, sulfur atoms can replace some oxygen atoms in carbon–oxygen–central inorganic atom covalent bonds of an ester, esters contain a carbonyl center, which gives rise to 120 ° C–C–O and O–C–O angles. Unlike amides, esters are structurally flexible functional groups because rotation about the C–O–C bonds has a low barrier and their flexibility and low polarity is manifested in their physical properties, they tend to be less rigid and more volatile than the corresponding amides. The pKa of the alpha-hydrogens on esters is around 25, the preference for the Z conformation is influenced by the nature of the substituents and solvent, if present
24.
Tetrahedral molecular geometry
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In a tetrahedral molecular geometry, a central atom is located at the center with four substituents that are located at the corners of a tetrahedron. The bond angles are cos−1 =109.4712206. ° ≈109. 5° when all four substituents are the same, the perfectly symmetrical tetrahedron belongs to point group Td, but most tetrahedral molecules have lower symmetry. Aside from virtually all saturated organic compounds, most compounds of Si, Ge, often tetrahedral molecules feature multiple bonding to the outer ligands, as in xenon tetroxide, the perchlorate ion, the sulfate ion, the phosphate ion. Thiazyl trifluoride is tetrahedral, featuring a triple bond. Other molecules have an arrangement of electron pairs around a central atom. However the usual classification considers only the atoms and not the lone pair. The H–N–H angles are 107°, contracted from 109.5 and this difference is attributed to the influence of the lone pair which exerts a greater repulsive influence than a bonded atom. Again the geometry is widespread, particularly so for complexes where the metal has d0 or d10 configuration, illustrative examples include tetrakispalladium, nickel carbonyl, and titanium tetrachloride. Many complexes with incompletely filled d-shells are often tetrahedral, e. g. the tetrahalides of iron, cobalt, however in liquid water or in ice, the lone pairs form hydrogen bonds with neighboring water molecules. The most common arrangement of hydrogen atoms around an oxygen is tetrahedral with two hydrogen atoms bonded to oxygen and two attached by hydrogen bonds. Since the hydrogen bonds vary in many of these water molecules are not symmetrical. Many compounds and complexes adopt bitetrahedral structures, in this motif, the two tetrahedra share a common edge. The inorganic polymer silicon disulfide features a chain of edge-shared tetrahedra. Inversion of tetrahedral occurs widely in organic and main group chemistry, the so-called Walden inversion illustrates the stereochemical consequences of inversion at carbon. Nitrogen inversion in ammonia also entails transient formation of planar NH3, geometrical constraints in a molecule can cause a severe distortion of idealized tetrahedral geometry. In compounds featuring inverted tetrahedral geometry at a carbon atom, all four groups attached to carbon are on one side of a plane. The carbon atom lies at or near the apex of a pyramid with the other four groups at the corners. The simplest examples of organic molecules displaying inverted tetrahedral geometry are the smallest propellanes, such as propellane, or more generally the paddlanes, such molecules are typically strained, resulting in increased reactivity
25.
VSEPR theory
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Valence shell electron pair repulsion theory is a model used in chemistry to predict the geometry of individual molecules from the number of electron pairs surrounding their central atoms. It is also named the Gillespie-Nyholm theory after its two main developers, the acronym VSEPR is pronounced either ves-pur or vuh-seh-per. Gillespie has emphasized that the repulsion due to the Pauli exclusion principle is more important in determining molecular geometry than the electrostatic repulsion. VSEPR theory is based on electron density rather than mathematical wave functions and hence unrelated to orbital hybridisation. While it is qualitative, VSEPR has a quantitative basis in quantum chemical topology methods such as the electron localization function. In 1957, Ronald Gillespie and Ronald Sydney Nyholm of University College London refined this concept into a detailed theory. In VSEPR theory, a bond or triple bond are treated as a single bonding group. The sum of the number of atoms bonded to a central atom, the number of electron pairs, therefore, determines the overall geometry that they will adopt. For example, when there are two electron pairs surrounding the atom, their mutual repulsion is minimal when they lie at opposite poles of the sphere. Therefore, the atom is predicted to adopt a linear geometry. If there are 3 electron pairs surrounding the atom, their repulsion is minimized by placing them at the vertices of an equilateral triangle centered on the atom. Therefore, the geometry is trigonal. Likewise, for 4 electron pairs, the arrangement is tetrahedral. The overall geometry is further refined by distinguishing between bonding and nonbonding electron pairs, the bonding electron pair shared in a sigma bond with an adjacent atom lies further from the central atom than a nonbonding pair of that atom, which is held close to its positively charged nucleus. VSEPR theory therefore views repulsion by the pair to be greater than the repulsion by a bonding pair. As such, when a molecule has 2 interactions with different degrees of repulsion, for instance, when 5 valence electron pairs surround a central atom, they adopt a trigonal bipyramidal molecular geometry with two collinear axial positions and three equatorial positions. The difference between pairs and bonding pairs may also be used to rationalize deviations from idealized geometries. For example, the H2O molecule has four pairs in its valence shell
26.
Gilbert N. Lewis
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Lewis successfully contributed to thermodynamics, photochemistry, and isotope separation, and is also known for his concept of acids and bases. G. N. Lewis was born in 1875 in Weymouth, several years later, he became the Dean of the college of Chemistry at Berkeley, where he spent the rest of his life. As a professor, he incorporated thermodynamic principles into the chemistry curriculum and he began measuring the free energy values related to several chemical processes, both organic and inorganic. In 1916, he proposed his theory of bonding and added information about electrons in the periodic table of the chemical elements. In 1933, he started his research on isotope separation, Lewis worked with hydrogen and managed to purify a sample of heavy water. He then came up with his theory of acids and bases, in 1926, Lewis coined the term photon for the smallest unit of radiant energy. He was a brother in Alpha Chi Sigma, the chemistry fraternity. Though he was nominated 41 times, G. N. Lewis never won the Nobel Prize in Chemistry. On March 23,1946, Lewis was found dead in his Berkeley laboratory where he had working with hydrogen cyanide. After Lewis death, his children followed their fathers career in chemistry, Lewis was born in 1875 and raised in Weymouth, Massachusetts, where there exists a street named for him, G. N. Additionally, the wing of the new Weymouth High School Chemistry department has named in his honor. Lewis received his education at home from his parents, Frank Wesley Lewis, a lawyer of independent character. He read at age three and was intellectually precocious, in 1884 his family moved to Lincoln, Nebraska, and in 1889 he received his first formal education at the university preparatory school. In 1893, after two years at the University of Nebraska, Lewis transferred to Harvard University, where he obtained his B. S. in 1896. After a year of teaching at Harvard, Lewis took a fellowship to Germany, the center of physical chemistry. While working in Nernsts lab, Nernst and Lewis apparently developed a lifelong enmity, in the following years, Lewis started to criticize and denounce his former teacher on many occasions, calling Nernsts work on his heat theorem a regrettable episode in the history of chemistry. A friend of Nernsts, Wilhelm Palmær, was a member of the Nobel Chemistry Committee, after his stay in Nernsts lab, Lewis returned to Harvard in 1901 as an instructor for three more years. He was appointed instructor in thermodynamics and electrochemistry, in 1904 Lewis was granted a leave of absence and became Superintendent of Weights and Measures for the Bureau of Science in Manila, Philippines
27.
Formal charge
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In chemistry, a formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity. When determining the best Lewis structure for a molecule, the structure is such that the formal charge on each of the atoms is as close to zero as possible. Example, CO2 is a molecule with 16 total valence electrons. The following is equivalent, Draw a circle around the atom for which the charge is requested Count up the number of electrons in the atoms circle. Since the circle cuts the covalent bond in half, each covalent bond counts as one instead of two. Subtract the number of electrons in the circle from the number of the element to determine the formal charge. The formal charges computed for the atoms in this Lewis structure of carbon dioxide are shown below. It is important to keep in mind that formal charges are just that – formal, the formal charge system is just a method to keep track of all of the valence electrons that each atom brings with it when the molecule is formed. Formal charge is a tool for estimating the distribution of charge within a molecule. The concept of oxidation states constitutes a competing method to assess the distribution of electrons in molecules, with formal charge, the electrons in each covalent bond are assumed to be split exactly evenly between the two atoms in the bond. This can be most effectively visualized in an electrostatic potential map, with the oxidation state formalism, the electrons in the bonds are awarded to the atom with the greater electronegativity. In reality, the distribution of electrons in the molecule lies somewhere between two extremes
28.
Linus Pauling
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Linus Carl Pauling was an American chemist, biochemist, peace activist, author, and educator. He published more than 1,200 papers and books, of which about 850 dealt with scientific topics, new Scientist called him one of the 20 greatest scientists of all time, and as of 2000, he was rated the 16th most important scientist in history. Pauling was one of the founders of the fields of quantum chemistry, for his scientific work, Pauling was awarded the Nobel Prize in Chemistry in 1954. In 1962, for his activism, he was awarded the Nobel Peace Prize. This makes him the person to be awarded two unshared Nobel Prizes. He is one of four individuals to have won more than one Nobel Prize. Pauling is also one of two people to be awarded Nobel Prizes in different fields, the other being Marie Curie. Pauling also worked on DNAs structure, a problem which was solved by James Watson, Francis Crick, Rosalind Franklin, in his later years he promoted nuclear disarmament, as well as orthomolecular medicine, megavitamin therapy, and dietary supplements. None of the latter have gained acceptance in the scientific community. Pauling was born in Portland, Oregon, the child of Herman Henry William Pauling. He was named Linus Carl, in honor of Lucys father, Linus, in 1902, after his sister Pauline was born, Paulings parents decided to move out of Portland, to find more affordable and spacious living quarters than their one-room apartment. Lucy stayed with her husbands parents in Lake Oswego until Herman brought the family to Salem, within a year of Luciles birth in 1904, Herman Pauling moved his family to Oswego, where he opened his own drugstore. He moved his family to Condon, Oregon in 1905, by 1906, Herman Pauling was suffering from recurrent abdominal pain. He died of an ulcer on June 11,1910, leaving Lucy to care for Linus, Lucile. Pauling attributes his interest in becoming a chemist to being amazed by experiments conducted by a friend, Lloyd A. Jeffress, in high school, Pauling conducted chemistry experiments by scavenging equipment and material from an abandoned steel plant. With an older friend, Lloyd Simon, Pauling set up Palmon Laboratories in Simons basement and they approached local dairies offering to perform butterfat samplings at cheap prices but dairymen were wary of trusting two boys with the task, and the business ended in failure. At age 15, the high school senior had enough credits to enter Oregon State University, lacking two American history courses required for his high school diploma, Pauling asked the school principal if he could take the courses concurrently during the spring semester. Denied, he left Washington High School in June without a diploma, the school awarded him a diploma 45 years later, after he had won two Nobel Prizes
29.
Valence bond theory
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In chemistry, valence bond theory is one of two basic theories, along with molecular orbital theory, that were developed to use the methods of quantum mechanics to explain chemical bonding. It focuses on how the orbitals of the dissociated atoms combine to give individual chemical bonds when a molecule is formed. In contrast, molecular orbital theory has orbitals that cover the whole molecule, in 1916, G. N. Lewis proposed that a chemical bond forms by the interaction of two shared bonding electrons, with the representation of molecules as Lewis structures. In 1927 the Heitler–London theory was formulated which for the first time enabled the calculation of bonding properties of the hydrogen molecule H2 based on quantum mechanical considerations. Specifically, Walter Heitler determined how to use Schrödingers wave equation to show how two hydrogen atom wavefunctions join together, with plus, minus, and exchange terms, to form a covalent bond. He then called up his associate Fritz London and they worked out the details of the theory over the course of the night. Later, Linus Pauling used the pair bonding ideas of Lewis together with Heitler–London theory to two other key concepts in VB theory, resonance and orbital hybridization. Resonance theory was criticized as imperfect by Soviet chemists during the 1950s, according to this theory a covalent bond is formed between the two atoms by the overlap of half filled valence atomic orbitals of each atom containing one unpaired electron. A valence bond structure is similar to a Lewis structure, but where a single Lewis structure cannot be written, each of these VB structures represents a specific Lewis structure. This combination of valence bond structures is the point of resonance theory. Valence bond theory considers that the atomic orbitals of the participating atoms form a chemical bond. Because of the overlapping, it is most probable that electrons should be in the bond region, valence bond theory views bonds as weakly coupled orbitals. Valence bond theory is typically easier to employ in ground state molecules, the inner-shell orbitals and electrons remain essentially unchanged during the formation of bonds. The overlapping atomic orbitals can differ, the two types of overlapping orbitals are sigma and pi. Sigma bonds occur when the orbitals of two shared electrons overlap head-to-head, pi bonds occur when two orbitals overlap when they are parallel. For example, a bond between two electrons is a sigma bond, because two spheres are always coaxial. In terms of order, single bonds have one sigma bond, double bonds consist of one sigma bond and one pi bond. However, the orbitals for bonding may be hybrids
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Resonance (chemistry)
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In chemistry, resonance or mesomerism is a way of describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis structure. A molecule or ion with such delocalized electrons is represented by several contributing structures, each contributing structure can be represented by a Lewis structure, with only an integer number of covalent bonds between each pair of atoms within the structure. Several Lewis structures are used collectively to describe the molecular structure. Electron delocalization lowers the energy of the substance and thus makes it more stable than any of the contributing structures. The difference between the energy of the actual structure and that of the contributing structure with the lowest potential energy is called the resonance energy or delocalization energy. An isomer is a molecule with the chemical formula but with different arrangements of atoms in space. Resonance contributors of a molecule, on the contrary, can differ by the arrangements of electrons. Therefore, the resonance hybrid cannot be represented by a combination of isomers, benzene undergoes substitution reactions, rather than addition reactions as typical for alkenes. He proposed that the bond in benzene is intermediate of a single and double bond. The mechanism of resonance was introduced into quantum mechanics by Werner Heisenberg in 1926 in a discussion of the states of the helium atom. He compared the structure of the atom with the classical system of resonating coupled harmonic oscillators. Linus Pauling used this mechanism to explain the partial valence of molecules in 1928, the alternative term mesomerism popular in German and French publications with the same meaning was introduced by C. K. Ingold in 1938, but did not catch on in the English literature. The current concept of effect has taken on a related. The double headed arrow was introduced by the German chemist Fritz Arndt who preferred the German phrase zwischenstufe or intermediate stage, the real structure is an intermediate of these structures represented by a resonance hybrid. The contributing structures are not isomers and they differ only in the position of electrons, not in the position of nuclei. Each Lewis formula must have the number of valence electrons. Bonds that have different bond orders in different contributing structures do not have typical bond lengths, the real structure has a lower total potential energy than each of the contributing structures would have. This means that it is more stable than each separate contributing structure would be and it is a common misconception that resonance structures are actual transient states of the molecule, with the molecule oscillating between them or existing as an equilibrium between them
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Pi bond
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In chemistry, pi bonds are covalent chemical bonds where two lobes of one involved atomic orbital overlap two lobes of the other involved atomic orbital. Each of these orbitals is zero at a shared nodal plane. The same plane is also a plane for the molecular orbital of the pi bond. The Greek letter π in their name refers to p orbitals, P orbitals often engage in this sort of bonding. D orbitals also engage in pi bonding, and form part of the basis for metal-metal multiple bonding, Pi bonds are usually weaker than sigma bonds. From the perspective of quantum mechanics, this weakness is explained by significantly less overlap between the component p-orbitals due to their parallel orientation. This is contrasted by sigma bonds which form bonding orbitals directly between the nuclei of the atoms, resulting in greater overlap and a strong sigma bond. Pi bonds result from overlap of atomic orbitals that are in contact through two areas of overlap, pi-bonds are more diffuse bonds than the sigma bonds. Electrons in pi bonds are referred to as pi electrons. Molecular fragments joined by a pi bond cannot rotate about that bond without breaking the pi bond, for homonuclear diatomic molecules, bonding π molecular orbitals have only the one nodal plane passing through the bonded atoms, and no nodal planes between the bonded atoms. The corresponding antibonding, or π* molecular orbital, is defined by the presence of a nodal plane between these two bonded atoms. A typical double bond consists of one bond and one pi bond, for example. A typical triple bond, for example in acetylene, consists of one sigma bond, two pi bonds are the maximum that can exist between a given pair of atoms. Quadruple bonds are rare and can be formed only between transition metal atoms, and consist of one sigma bond, two pi bonds and one delta bond. A pi bond is weaker than a bond, but the combination of pi. The enhanced strength of a multiple bond versus a single is indicated in many ways, for example, in organic chemistry, carbon–carbon bond lengths are about 154 pm in ethane,134 pm in ethylene and 120 pm in acetylene. More bonds make the total bond shorter and stronger, a pi bond can exist between two atoms that do not have a net sigma-bonding effect between them. In certain metal complexes, pi interactions between an atom and alkyne and alkene pi antibonding orbitals form pi-bonds
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Pauling's principle of electroneutrality
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Paulings principle of electroneutrality states that each atom in a stable substance has a charge close to zero. It was formulated by Linus Pauling in 1948 and later revised, however, modern computational techniques indicate many stable compounds have a greater charge distribution than the principle predicts. Pauling first stated his postulate of the essential electroneutrality of atoms in 1948, at the Liversidge lecture, a slightly revised version was published in 1970, “Stable molecules and crystals have electronic structures such that the electric charge of each atom is close to zero. Close to zero means between -1 and +1. ”Pauling said in his Liversidge lecture in 1948 that he had led to the principle by a consideration of ionic bonding. In the gas phase, molecular caesium fluoride has a covalent bond. The large difference in electronegativity gives a calculated ionic character of 9%, in the crystal if each bond has 9% covalent character the total covalency of Cs and F would be 54%. This would be represented by one bond resonating between the six positions and the effect would be to reduce the charge on Cs to about +0.5. It seemed reasonable to him such a reduction would be general. There are two possible structures for hydrogen cyanide, HCN and CNH, differing only as to the position of the hydrogen atom.79, +0.75 and +0.04 respectively.04, +0.17, the triple bonded structure is therefore favored. Applying the principle of electroneutrality this can be identified as only a minor contributor, additionally as the most electronegative atom should carry the negative charge, then the triple bonded structure on the left is predicted to be the major contributor. The hexaammine cobalt complex 3+ would have all of charge on the central Co atom if the bonding to the molecules were electrostatic. On the other hand, a covalent linkage would lead to a charge of -3 on the metal, using the electroneutrality principle the assumption is made that the Co-N bond will have 50% ionic character thus resulting in a zero charge on the cobalt atom. Due to the difference in electronegativity the N-H bond would 17% ionic character and this essentially spreads the 3+ charge evenly onto the surface of the complex ion. Pauling invoked the principle of electroneutrality in a 1952 paper to suggest that pi bonding is present, for example, the oxygen atoms in such molecules would form polar covalent bonds with the silicon atom because their electronegativity was higher than that of silicon. Pauling calculated the charge build up on the silicon atom due to the difference in electronegativity to be +2, the adjacent charge rule was another principle of Paulings for determining whether a resonance structure would make a significant contribution. First published in 1932, it stated that structures that placed charges of the sign on adjacent atoms would be unfavorable
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Natural bond orbital
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In quantum chemistry, a natural bond orbital or NBO is a calculated bonding orbital with maximum electron density. The NBOs are one of a sequence of natural localized orbital sets that include natural atomic orbitals, natural hybrid orbitals, natural bonding orbitals and they have the maximum-occupancy character in localized 1-center and 2-center regions of the molecule. Natural bond orbitals include the highest possible percentage of the density, ideally close to 2.000. A high percentage of electron density, often found to be >99% for common organic molecules, the concept of natural orbitals was first introduced by Per-Olov Löwdin in 1955, to describe the unique set of orthonormal 1-electron functions that are intrinsic to the N-electron wavefunction. In an idealized Lewis structure, full Lewis orbitals are complemented by formally empty non-Lewis orbitals, weak occupancies of the valence antibonds signal irreducible departures from an idealized localized Lewis structure, which means true delocalization effects. With a computer program that can calculate NBOs, optimal Lewis structures can be found, an optimal Lewis structure can be defined as that one with the maximum amount of electronic charge in Lewis orbitals. A low amount of charge in Lewis orbitals indicates strong effects of electron delocalization. In resonance structures, major and minor contributing structures may exist, for amides, for example, NBO calculations show that the structure with a carbonyl double bond is the dominant Lewis structure. However, in NBO calculations, covalent-ionic resonance is not needed due to the inclusion of bond-polarity effects in the resonance structures and this is similar to other modern valence bond theory methods
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Octet rule
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The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens, but also to metals such as sodium or magnesium. The valence electrons can be counted using a Lewis electron dot diagram as shown at the right for carbon dioxide, the electrons shared by the two atoms in a covalent bond are counted twice, once for each atom. In carbon dioxide each oxygen shares four electrons with the central carbon, all four of these electrons are counted in both the carbon octet and the oxygen octet. Ionic bonding is common between pairs of atoms, where one of the pair is a metal of low electronegativity, a chlorine atom has seven electrons in its outer electron shell, the first and second shells being filled with two and eight electrons respectively. The first electron affinity of chlorine is +328.8 kJ per mole of chlorine atoms, adding a second electron to chlorine requires energy, energy that cannot be recovered by formation of a chemical bond. The result is that chlorine will very often form a compound in which it has eight electrons in its outer shell, a sodium atom has a single electron in its outermost electron shell, the first and second shells again being full with two and eight electrons respectively. To remove this outer electron requires only the first ionization energy, which is +495.8 kJ per mole of sodium atoms, a small amount of energy. By contrast, the electron resides in the deeper second electron shell. Thus sodium will, in most cases, form a compound in which it has lost an electron and have a full outer shell of eight electrons. The energy required to transfer an electron from an atom to a chlorine atom is small. This energy is offset by the lattice energy of sodium chloride. This completes the explanation of the rule in this case. In 1893, Alfred Werner showed that the number of atoms or groups associated with an atom is often 4 or 6, other coordination numbers up to a maximum of 8 were known. Abegg noted that the difference between the positive and negative valences of an element under his model is frequently eight. In 1919 Irving Langmuir refined these concepts further and renamed them the cubical octet atom, the octet theory evolved into what is now known as the octet rule. The quantum theory of the atom explains the eight electrons as a shell with an s2p6 electron configuration. A closed-shell configuration is one in which low-lying energy levels are full, for example, the neon atom ground state has a full n =2 shell and an empty n =3 shell. According to the rule, the atoms immediately before and after neon in the periodic table, tend to attain a similar configuration by gaining, losing
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Electronegativity
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Electronegativity, symbol χ, is a chemical property that describes the tendency of an atom to attract electrons towards itself. An atoms electronegativity is affected by both its number and the distance at which its valence electrons reside from the charged nucleus. The higher the associated electronegativity number, the more an element or compound attracts electrons towards it, the term electronegativity was introduced by Jöns Jacob Berzelius in 1811, though the concept was known even before that and was studied by many chemists including Avogadro. It has been shown to correlate with a number of chemical properties. Electronegativity cannot be measured and must be calculated from other atomic or molecular properties. The most commonly used method of calculation is that proposed by Linus Pauling. This gives a quantity, commonly referred to as the Pauling scale. When other methods of calculation are used, it is conventional to quote the results on a scale that covers the range of numerical values. As it is calculated, electronegativity is not a property of an atom alone. Properties of a free atom include ionization energy and electron affinity, on the most basic level, electronegativity is determined by factors like the nuclear charge and the number/location of other electrons present in the atomic shells. The opposite of electronegativity is electropositivity, a measure of an ability to donate electrons. Caesium is the least electronegative element in the table, while fluorine is most electronegative. According to valence bond theory, of which Pauling was a notable proponent, hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements, its electronegativity was fixed first at 2.1, later revised to 2.20. It is also necessary to decide which of the two elements is the more electronegative, to calculate Pauling electronegativity for an element, it is necessary to have data on the dissociation energies of at least two types of covalent bond formed by that element. A. L.32 e V This is an approximate equation, Pauling obtained it by noting that a bond can be approximately represented as a quantum mechanical superposition of a covalent bond and two ionic bond-states. e. The geometric mean is equal to the arithmetic mean - which is applied in the first formula above - when the energies are of the similar value. The square root of this energy, Pauling notes, is approximately additive. Thus, it is this semi-empirical formula for energy that underlies Pauling electronegativity concept
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Covalent bond
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A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, for many molecules, the sharing of electrons allows each atom to attain the equivalent of a full outer shell, corresponding to a stable electronic configuration. Covalent bonding includes many kinds of interactions, including σ-bonding, π-bonding, metal-to-metal bonding, agostic interactions, bent bonds, the term covalent bond dates from 1939. In the molecule H2, the atoms share the two electrons via covalent bonding. Covalency is greatest between atoms of similar electronegativities, thus, covalent bonding does not necessarily require that the two atoms be of the same elements, only that they be of comparable electronegativity. Covalent bonding that entails sharing of electrons more than two atoms is said to be delocalized. The term covalence in regard to bonding was first used in 1919 by Irving Langmuir in a Journal of the American Chemical Society article entitled The Arrangement of Electrons in Atoms and Molecules. Langmuir wrote that we shall denote by the term covalence the number of pairs of electrons that an atom shares with its neighbors. The idea of covalent bonding can be traced several years before 1919 to Gilbert N. Lewis and he introduced the Lewis notation or electron dot notation or Lewis dot structure, in which valence electrons are represented as dots around the atomic symbols. Pairs of electrons located 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 bond-forming electron pairs represented as solid lines, Lewis proposed that an atom forms enough covalent bonds to form a full outer electron shell. In the diagram of methane shown here, the atom has a valence of four and is, therefore, surrounded by eight electrons, four from the carbon itself. Each hydrogen has a valence of one and is surrounded by two electrons – its own one electron plus one from the carbon, walter Heitler and Fritz London are credited with the first successful quantum mechanical explanation of a chemical bond in 1927. Their work was based on the valence bond model, which assumes that a bond is formed when there is good overlap between the atomic orbitals of participating atoms. Atomic orbitals have specific directional properties leading to different types of covalent bonds, sigma bonds are the strongest covalent bonds and are due to head-on overlapping of orbitals on two different atoms. A single bond is usually a σ bond, pi bonds are weaker and are due to lateral overlap between p orbitals. A double bond between two given atoms consists of one σ and one π bond, and a bond is one σ. Covalent bonds are also affected by the electronegativity of the atoms which determines the chemical polarity of the bond
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Ionic bonding
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Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, and is the primary interaction occurring in ionic compounds. The ions are atoms that have gained one or more electrons and this transfer of electrons is known as electrovalence in contrast to covalence. In the simplest case, the cation is an atom and the anion is a nonmetal atom. In simpler words, a bond is the transfer of electrons from a metal to a non-metal in order for both atoms to obtain a full valence shell. Bonds with partially ionic and partially covalent character are called polar covalent bonds, Ionic compounds conduct electricity when molten or in solution, typically as a solid. Ionic compounds generally have a melting point, depending on the charge of the ions they consist of. The higher the charges the stronger the forces and the higher the melting point. They also tend to be soluble in water, here, the opposite trend roughly holds, the weaker the cohesive forces, the greater the solubility. Atoms that have an almost full or almost empty valence shell tend to be very reactive, atoms that are strongly electronegative often have only one or two empty orbitals in their valence shell, and frequently bond with other molecules or gain electrons to form anions. Atoms that are weakly electronegative have relatively few valence electrons that can easily be shared with atoms that are strongly electronegative, as a result, weakly electronegative atoms tend to distort their electrons cloud and form cations. Ionic bonding can result from a reaction when atoms of an element, whose ionization energy is low. In doing so, cations are formed, the atom of another element, whose electron affinity is positive, then accepts the electron, again to attain a stable electron configuration, and after accepting electron the atom becomes an anion. Typically, the electron configuration is one of the noble gases for elements in the s-block and the p-block. The electrostatic attraction between the anions and cations leads to the formation of a solid with a lattice in which the ions are stacked in an alternating fashion. In such a lattice, it is not possible to distinguish discrete molecular units. However, the ions themselves can be complex and form molecular ions like the acetate anion or the ammonium cation, for example, common table salt is sodium chloride. When sodium and chlorine are combined, the sodium atoms each lose an electron, forming cations, and these ions are then attracted to each other in a 1,1 ratio to form sodium chloride. For compounds that are transitional to the alloys and possess mixed ionic and metallic bonding, many sulfides, e. g. do form non-stoichiometric compounds
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Lone pair
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In chemistry, a lone pair refers to a pair of valence electrons that are not shared with another atom and is sometimes called a non-bonding pair. Lone pairs are found in the outermost electron shell of atoms and they can be identified by using a Lewis structure. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding, thus, the number of lone pair electrons plus the number of bonding electrons equals the total number of valence electrons around an atom. Lone pair is a used in VSEPR theory which explains the shapes of molecules. They are also referred to in the chemistry of Lewis acids, however, not all non-bonding pairs of electrons are considered by chemists to be lone pairs. Examples are the transition metals where the non-bonding pairs do not influence molecular geometry and are said to be stereochemically inactive, in VSEPR theory the electron pairs on the oxygen atom in water form the vertices of a tetrahedron with the lone pairs on two of the four vertices. The H–O–H bond angle is with 104. 5°, less than the 109° predicted for an angle. The pairs often exhibit a negative polar character with their charge density and are located closer to the atomic nucleus on average compared to the bonding pair of electrons. The presence of a lone pair decreases the angle between the bonding pair of electrons, due to their high electric charge which causes great repulsion between the electrons. They are also used in the formation of a dative bond, for example, the creation of the hydronium ion occurs when acids are dissolved in water and is due to the oxygen atom donating a lone pair to the hydrogen ion. This can be more clearly when looked at it in two more common molecules. For example, in carbon dioxide, the atoms are on opposite sides of the carbon. Due to the force of the oxygen atoms lone pairs. This is an illustration of the VSEPR theory, lone pairs can make a contribution to a molecules dipole moment. NH3 has a moment of 1.47 D. There is also an associated with the lone pair and this reinforces the contribution made by the polar covalent N-H bonds to ammonias dipole moment. In contrast to NH3, NF3 has a lower dipole moment of 0.24 D. A stereochemically active lone pair is expected for divalent lead
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Sulfide
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Sulfide is an inorganic anion of sulfur with the chemical formula S2−. It contributes no color to sulfide salts, as it is classified as a strong base, even dilute solutions of salts such as sodium sulfide are corrosive and can attack the skin. Sulfide is the simplest sulfur anion, the systematic names sulfanediide and sulfide, valid IUPAC names, are determined according to the substitutive and additive nomenclatures, respectively. However, the sulfide is also used in compositional IUPAC nomenclature which does not take the nature of bonding involved. Examples of such naming are selenium disulfide and titanium sulfide, which contains no sulfide ions whatsoever, hydrogen sulfide is itself an example of a non-systematic name of this nature. However, it is also a name, and the preferred IUPAC name for sulfane. Sulfide does not exist in appreciable concentrations even in highly alkaline water, the sulfide anion can assimilate a proton by recombination, S2− + H+ → SH− Because of this capture of a proton, sulfide has basic character. In aqueous solution, it has a pKb value of less than 0, in aqueous solution, most sulfide ions are neutralized. S2− + H2O ↽ − ⇀ SH− + OH− Upon treatment with an acid, sulfide converts to hydrogen sulfide. Oxidation of sulfide gives sulfur or sulfate, metal sulfides react with nonmetals including iodine, bromine, and chlorine forming sulfur and metal salts. Such inorganic sulfides typically have low solubility in water. One famous example is the bright yellow species CdS or cadmium yellow, the black tarnish formed on sterling silver is Ag2S. Such species are referred to as salts. In fact, the bonding in transition metal sulfides is highly covalent, which rise to their semiconductor properties. Several have practical applications as pigments, in cells. Aspergillus niger plays a role in the solubilization of heavy metal sulfides, many important metal ores are sulfides. Significant examples include, argentite, cinnabar, galena, molybdenite, pentlandite, realgar, and stibnite, sphalerite, and pyrite, dissolved free sulfides are very aggressive species for the corrosion of many metals such as steel, stainless steel, and copper. Sulfides present in solution are responsible for stress corrosion cracking of steel
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Sulfite
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Sulfites or sulphites are compounds that contain the sulfite ion, SO2−3. The sulfite ion is the base of bisulfite. Although its acid is elusive, its salts are widely used, Sulfites are substances that naturally occur in some foods and the human body. They are also used as regulated food additives, the structure of the sulfite anion can be described with three equivalent resonance structures. There is also a lone pair on the sulfur, so the structure predicted by VSEPR theory is trigonal pyramidal. In the hybrid resonance structure, the S-O bonds are equivalently of bond order one, evidence from 17O NMR spectroscopic data suggests that protonation of the sulfite ion gives a mixture of isomers, Sulfites are used as a food preservative or enhancer. Sulfites are commonly introduced to arrest fermentation at a desired time, sulfur dioxide protects wine from not only oxidation, but also from bacteria. Without sulfites, grape juice would turn to vinegar. Organic wines are not necessarily sulfite-free, but generally have lower amounts, in general, white wines contain more sulfites than red wines and sweeter wines contain more sulfites than drier ones. In the United States, wines bottled after mid-1987 must have a label stating that they contain sulfites if they contain more than 10 parts per million, in the European Union an equivalent regulation came into force in November 2005. In 2012, a new regulation for organic wines came into force, Sulfites are often used as preservatives in dried fruits, preserved radish, and dried potato products. Most beers no longer contain sulfites, although some alcoholic ciders contain them, although shrimp are sometimes treated with sulfites on fishing vessels, the chemical may not appear on the label. In 1986, the Food and Drug Administration in the United States banned the addition of sulfites to all fresh fruit, E numbers for sulfites as food additives are, Sulfites are counted among the top nine food allergens, but a reaction to sulfite is not a true allergy. Some people have positive skin allergy tests to sulfites indicating true allergy and it may cause breathing difficulty within minutes after eating a food containing it. Asthmatics and possibly people with sensitivity are at an elevated risk for reaction to sulfites. Anaphylaxis and life-threatening reactions are rare, other potential symptoms include sneezing, swelling of the throat, hives, and migraine. In 1986, the U. S. Food and Drug Administration banned the use of sulfites as preservatives on foods intended to be eaten fresh and this has contributed to the increased use of erythorbic acid and its salts as preservatives. In the U. S. Sulfites used in food processing, the products most likely to contain sulfites do not require ingredients labels, so the presence of sulfites usually is undisclosed
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Solubility
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Solubility is the property of a solid, liquid, or gaseous chemical substance called solute to dissolve in a solid, liquid, or gaseous solvent. The solubility of a substance depends on the physical and chemical properties of the solute and solvent as well as on temperature, pressure. The solubility of a substance is a different property from the rate of solution. Most often, the solvent is a liquid, which can be a substance or a mixture. One may also speak of solid solution, but rarely of solution in a gas, the extent of solubility ranges widely, from infinitely soluble such as ethanol in water, to poorly soluble, such as silver chloride in water. The term insoluble is often applied to poorly or very poorly soluble compounds, a common threshold to describe something as insoluble is less than 0.1 g per 100 mL of solvent. Under certain conditions, the solubility can be exceeded to give a so-called supersaturated solution. Metastability of crystals can also lead to apparent differences in the amount of a chemical that dissolves depending on its form or particle size. A supersaturated solution generally crystallises when seed crystals are introduced and rapid equilibration occurs, phenylsalicylate is one such simple observable substance when fully melted and then cooled below its fusion point. Solubility is not to be confused with the ability to dissolve a substance, for example, zinc dissolves in hydrochloric acid as a result of a chemical reaction releasing hydrogen gas in a displacement reaction. The zinc ions are soluble in the acid, the smaller a particle is, the faster it dissolves although there are many factors to add to this generalization. Crucially solubility applies to all areas of chemistry, geochemistry, inorganic, physical, organic, in all cases it will depend on the physical conditions and the enthalpy and entropy directly relating to the solvents and solutes concerned. By far the most common solvent in chemistry is water which is a solvent for most ionic compounds as well as a range of organic substances. This is a factor in acidity/alkalinity and much environmental and geochemical work. According to the IUPAC definition, solubility is the composition of a saturated solution expressed as a proportion of a designated solute in a designated solvent. Solubility may be stated in units of concentration such as molarity, molality, mole fraction, mole ratio, mass per volume. Solubility occurs under dynamic equilibrium, which means that solubility results from the simultaneous and opposing processes of dissolution, the solubility equilibrium occurs when the two processes proceed at a constant rate. The term solubility is used in some fields where the solute is altered by solvolysis
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