Orbital hybridisation

In chemistry, orbital hybridisation (or hybridization) is the concept of mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory. Hybrid orbitals are very useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. Although sometimes taught together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and hybridisation are in fact not related to the VSEPR model.[1]

History and uses

Chemist Linus Pauling first developed the hybridisation theory in 1931 to explain the structure of simple molecules such as methane (CH4) using atomic orbitals.[2] Pauling pointed out that a carbon atom forms four bonds by using one s and three p orbitals, so that "it might be inferred" that a carbon atom would form three bonds at right angles (using p orbitals) and a fourth weaker bond using the s orbital in some arbitrary direction. In reality, methane has four bonds of equivalent strength separated by the tetrahedral bond angle of 109.5°. Pauling explained this by supposing that in the presence of four hydrogen atoms, the s and p orbitals form four equivalent combinations or hybrid orbitals, each denoted by sp3 to indicate its composition, which are directed along the four C-H bonds.[3] This concept was developed for such simple chemical systems, but the approach was later applied more widely, and today it is considered an effective heuristic for rationalising the structures of organic compounds. It gives a simple orbital picture equivalent to Lewis structures.

Hybridisation theory is an integral part of organic chemistry, one of the most compelling examples being Baldwin's rules. For drawing reaction mechanisms sometimes a classical bonding picture is needed with two atoms sharing two electrons.[4] Hybridisation theory explains bonding in alkenes[5] and methane.[6] The amount of p character or s character, which is decided mainly by orbital hybridisation, can be used to reliably predict molecular properties such as acidity or basicity.[7]

Overview

Orbitals are a model representation of the behaviour of electrons within molecules. In the case of simple hybridisation, this approximation is based on atomic orbitals, similar to those obtained for the hydrogen atom, the only neutral atom for which the Schrödinger equation can be solved exactly. In heavier atoms, such as carbon, nitrogen, and oxygen, the atomic orbitals used are the 2s and 2p orbitals, similar to excited state orbitals for hydrogen.

Hybrid orbitals are assumed to be mixtures of atomic orbitals, superimposed on each other in various proportions. For example, in methane, the C hybrid orbital which forms each carbonhydrogen bond consists of 25% s character and 75% p character and is thus described as sp3 (read as s-p-three) hybridised. Quantum mechanics describes this hybrid as an sp3 wavefunction of the form N(s + 3pσ), where N is a normalisation constant (here 1/2) and pσ is a p orbital directed along the C-H axis to form a sigma bond. The ratio of coefficients (denoted λ in general) is 3 in this example. Since the electron density associated with an orbital is proportional to the square of the wavefunction, the ratio of p-character to s-character is λ2 = 3. The p character or the weight of the p component is N2λ2 = 3/4.

Types of hybridisation

sp3

AE4h
Four sp3 orbitals.

Hybridisation describes the bonding atoms from an atom's point of view. For a tetrahedrally coordinated carbon (e.g., methane CH4), the carbon should have 4 orbitals with the correct symmetry to bond to the 4 hydrogen atoms.

Carbon's ground state configuration is 1s2 2s2 2p2 or more easily read:

C ↑↓ ↑↓  
1s 2s 2p 2p 2p

The carbon atom can use its two singly occupied p-type orbitals, to form two covalent bonds with two hydrogen atoms, yielding the singlet methylene CH2, the simplest carbene. The carbon atom can also bond to four hydrogen atoms by an excitation (or promotion) of an electron from the doubly occupied 2s orbital to the empty 2p orbital, producing four singly occupied orbitals.

C* ↑↓
1s 2s 2p 2p 2p

The energy released by formation of two additional bonds more than compensates for the excitation energy required, energetically favouring the formation of four C-H bonds.

Quantum mechanically, the lowest energy is obtained if the four bonds are equivalent, which requires that they are formed from equivalent orbitals on the carbon. A set of four equivalent orbitals can be obtained that are linear combinations of the valence-shell (core orbitals are almost never involved in bonding) s and p wave functions,[8] which are the four sp3 hybrids.

C* ↑↓
1s sp3 sp3 sp3 sp3

In CH4, four sp3 hybrid orbitals are overlapped by hydrogen 1s orbitals, yielding four σ (sigma) bonds (that is, four single covalent bonds) of equal length and strength.

A schematic presentation of hybrid orbitals overlapping hydrogen orbitals
translates into
Methane's tetrahedral shape

A schematic presentation of hybrid orbitals overlapping hydrogen orbitals
Methane's tetrahedral shape

sp2

AE3h
Three sp2 orbitals.
Ethene-2D-flat
Ethene structure

Other carbon compounds and other molecules may be explained in a similar way. For example, ethene (C2H4) has a double bond between the carbons.

For this molecule, carbon sp2 hybridises, because one π (pi) bond is required for the double bond between the carbons and only three σ bonds are formed per carbon atom. In sp2 hybridisation the 2s orbital is mixed with only two of the three available 2p orbitals,

C* ↑↓
1s sp2 sp2 sp2 2p

forming a total of three sp2 orbitals with one remaining p orbital. In ethylene (ethene) the two carbon atoms form a σ bond by overlapping one sp2 orbital from each carbon atom. The π bond between the carbon atoms perpendicular to the molecular plane is formed by 2p–2p overlap. Each carbon atom forms covalent C–H bonds with two hydrogens by s–sp2 overlap, all with 120° bond angles. The hydrogen–carbon bonds are all of equal strength and length, in agreement with experimental data.

sp

AE2h
Two sp orbitals

The chemical bonding in compounds such as alkynes with triple bonds is explained by sp hybridisation. In this model, the 2s orbital is mixed with only one of the three p orbitals,

C* ↑↓
1s sp sp 2p 2p

resulting in two sp orbitals and two remaining p orbitals. The chemical bonding in acetylene (ethyne) (C2H2) consists of sp–sp overlap between the two carbon atoms forming a σ bond and two additional π bonds formed by p–p overlap. Each carbon also bonds to hydrogen in a σ s–sp overlap at 180° angles.

Hybridisation and molecule shape

Hybridisation helps to explain molecule shape, since the angles between bonds are (approximately) equal to the angles between hybrid orbitals, as explained above for the tetrahedral geometry of methane. As another example, the three sp2 hybrid orbitals are at angles of 120° to each other, so this hybridisation favours trigonal planar molecular geometry with bond angles of 120°. Other examples are given in the table below.

Classification Main group Transition metal[9][10]
AX2
  • Linear (180°)
  • sp hybridisation
  • E.g., CO2
  • Bent (90°)
  • sd hybridisation
  • E.g., VO2+
AX3
AX4
AX6

For two equivalent spx hybrids, the bond angle between them is given by , while for two equivalent sdx hybrids, the bond angle between them is given by .[9] There are no shapes with ideal bond angles corresponding to sd4 hybridisation as there is no regular 10-vertex polyhedron. Nonetheless, molecules such as Ta(CH3)5 with sd4 hybridisation adopt a square pyramidal shape.

Hybridisation of hypervalent molecules

Valence shell expansion

Hybridisation is often presented for main group AX5 and above, as well as for many transition metal complexes, using the hybridisation scheme first proposed by Pauling.

Classification Main group
AX5
AX6
AX7
Classification Transition metal[11]
AX4 (d8)
AX6 (d6)
  • Octahedral (90°)
  • d2sp3 hybridisation
  • E.g., Mo(CO)6
Note: AX5 and AX7 have variable molecular shapes[12][13]

In this notation, d orbitals of main group atoms are listed after the s and p orbitals since they have the same principal quantum number (n), while d orbitals of transition metals are listed first since the s and p orbitals have a higher n. Thus for AX6 molecules, sp3d2 hybridisation in the S atom involves 3s, 3p and 3d orbitals, while d2sp3 for Mo involves 4d, 5s and 5p orbitals.

Contrary evidence

In 1990, Magnusson published a seminal work definitively excluding the role of d-orbital hybridisation in bonding in hypervalent compounds of second-row (period 3) elements, ending a point of contention and confusion. Part of the confusion originates from the fact that d-functions are essential in the basis sets used to describe these compounds (or else unreasonably high energies and distorted geometries result). Also, the contribution of the d-function to the molecular wavefunction is large. These facts were incorrectly interpreted to mean that d-orbitals must be involved in bonding.[14][15]

For transition metal centres, the d and s orbitals are the primary valence orbitals, which are only weakly supplemented by the p orbitals.[16] The question of whether the p orbitals actually participate in bonding has not been definitively resolved, but all studies indicate they play a minor role.

Resonance

In light of computational chemistry, a better treatment would be to invoke sigma bond resonance in addition to hybridisation, which implies that each resonance structure has its own hybridisation scheme. For main group compounds, all resonance structures must obey the octet (8-electron) rule. For transition metal compounds, the resonance structures that obey the duodectet (12-electron) rule[17] suffice to describe bonding, with optional inclusion of dmspn resonance structures.

Classification Main group[18]
AX5 Trigonal bipyramidal (90°, 120°)
Penta phos
AX6 Octahedral (90°)
Hexa sulf
AX7 Pentagonal bipyramidal (90°, 72°)
Hepta iodi
Classification Transition metal[9]
AX4 (d8) Square planar (90°)
Tetra plat
AX6 (d6) Octahedral (90°)
Hexa moly
Note: AX5 and AX7 have variable molecular shapes[12][13]

Isovalent hybridisation

Although ideal hybrid orbitals can be useful, in reality most bonds require orbitals of intermediate character. This requires an extension to include flexible weightings of atomic orbitals of each type (s, p, d) and allows for a quantitative depiction of bond formation when the molecular geometry deviates from ideal bond angles. The amount of p-character is not restricted to integer values; i.e., hybridisations like sp2.5 are also readily described.

The hybridisation of bond orbitals is determined by Bent's rule: "Atomic s character concentrates in orbitals directed towards electropositive substituents".

Molecules with lone pairs

For molecules with lone pairs, the bonding orbitals are isovalent spx hybrids. For example, the two bond-forming hybrid orbitals of oxygen in water can be described as sp4.0 to give the interorbital angle of 104.5°.[19] This means that they have 20% s character and 80% p character and does not imply that a hybrid orbital is formed from one s and four p orbitals on oxygen since the 2p subshell of oxygen only contains three p orbitals. The shapes of molecules with lone pairs are:

In such cases, there are two mathematically equivalent ways of representing lone pairs. They can be represented by orbitals of sigma and pi symmetry similar to molecular orbital theory or by equivalent orbitals similar to VSEPR theory.

Hypervalent molecules

For hypervalent molecules with lone pairs, the bonding scheme can be split into a hypervalent component and a component consisting of isovalent spx bond hybrids. The hypervalent component consists of resonant bonds using p orbitals. The table below shows how each shape is related to the two components and their respective descriptions.

Number of spx bond hybrids (marked in red)
Two One
Hypervalent component[18] Linear axis
(one p orbital)
Seesaw (90°, 180°, >90°) T-shaped (90°, 180°) Linear (180°)
Tetra sulf Tri chlo Di xeno
Square planar equator
(two p orbitals)
Square pyramidal (90°, 90°) Square planar (90°)
Penta chlo Tetra xeno
Pentagonal planar equator
(two p orbitals)
Pentagonal pyramidal (90°, 72°) Pentagonal planar (72°)
Hexa xeno Penta xeno

Hybridisation defects

Hybridisation of s and p orbitals to form effective spx hybrids requires that they have comparable radial extent. While 2p orbitals are on average less than 10% larger than 2s, in part attributable to the lack of a radial node in 2p orbitals, 3p orbitals which have one radial node, exceed the 3s orbitals by 20–33%.[20] The difference in extent of s and p orbitals increases further down a group. The hybridisation of atoms in chemical bonds can be analysed by considering localised molecular orbitals, for example using natural localised molecular orbitals in a natural bond orbital (NBO) scheme. In methane, CH4, the calculated p/s ratio is approximately 3 consistent with "ideal" sp3 hybridisation, whereas for silane, SiH4, the p/s ratio is closer to 2. A similar trend is seen for the other 2p elements. Substitution of fluorine for hydrogen further decreases the p/s ratio.[21] The 2p elements exhibit near ideal hybridisation with orthogonal hybrid orbitals. For heavier p block elements this assumption of orthogonality cannot be justified. These deviations from the ideal hybridisation were termed hybridisation defects by Kutzelnigg.[22]

Photoelectron spectra

One misconception concerning orbital hybridisation is that it incorrectly predicts the ultraviolet photoelectron spectra of many molecules. While this is true if Koopmans' theorem is applied to localised hybrids, quantum mechanics requires that the (in this case ionised) wavefunction obey the symmetry of the molecule which implies resonance in valence bond theory. For example, in methane, the ionised states (CH4+) can be constructed out of four resonance structures attributing the ejected electron to each of the four sp3 orbitals. A linear combination of these four structures, conserving the number of structures, leads to a triply degenerate T2 state and a A1 state.[23] The difference in energy between each ionised state and the ground state would be an ionisation energy, which yields two values in agreement with experiment.

Localized MOs vs canonical MOs

Bonding orbitals formed from hybrid atomic orbitals may be considered as localized molecular orbitals, which can be formed from the delocalised orbitals of molecular orbital theory by an appropriate mathematical transformation. For molecules in the ground state, this transformation of the orbitals leaves the total many-electron wave function unchanged. The hybrid orbital description of the ground state is therefore equivalent to the delocalised orbital description for ground state total energy and electron density, as well as the molecular geometry that corresponds to the minimum total energy value.

Two localized representations

Molecules with multiple bonds or multiple lone pairs can have orbitals represented in terms of sigma and pi symmetry or equivalent orbitals. Different valence bond methods use either of the two representations, which have mathematically equivalent total many-electron wave functions and are related by a unitary transformation of the set of occupied molecular orbitals.

For multiple bonds, the sigma-pi representation is the predominant one compared to the equivalent orbital (bent bond) representation. In contrast, for multiple lone pairs, most textbooks use the equivalent orbital representation. However, the sigma-pi representation is also used, such as by Weinhold and Landis within the context of natural bond orbitals, a localized orbital theory containing modernized analogs of classical (valence bond/Lewis structure) bonding pairs and lone pairs.[24] For the hydrogen fluoride molecule, for example, two F lone pairs are essentially unhybridized p orbitals, while the other is an spx hydrid orbital. An analogous consideration applies to water (one O lone pair is in a pure p orbital, another is in an spx hybrid orbital).

See also

References

  1. ^ Gillespie, R.J. (2004), "Teaching molecular geometry with the VSEPR model", Journal of Chemical Education, 81 (3): 298–304, Bibcode:2004JChEd..81..298G, doi:10.1021/ed081p298
  2. ^ Pauling, L. (1931), "The nature of the chemical bond. Application of results obtained from the quantum mechanics and from a theory of paramagnetic susceptibility to the structure of molecules", Journal of the American Chemical Society, 53 (4): 1367–1400, doi:10.1021/ja01355a027
  3. ^ L. Pauling The Nature of the Chemical Bond (3rd ed., Oxford University Press 1960) p.111–120.
  4. ^ Clayden, Jonathan; Greeves, Nick; Warren, Stuart; Wothers, Peter (2001). Organic Chemistry (1st ed.). Oxford University Press. p. 105. ISBN 978-0-19-850346-0.
  5. ^ Organic Chemistry, Third Edition Marye Anne Fox James K. Whitesell 2003 ISBN 978-0-7637-3586-9
  6. ^ Organic Chemistry 3rd Ed. 2001 Paula Yurkanis Bruice ISBN 978-0-130-17858-9
  7. ^ "Acids and Bases". Orgo Made Simple. Retrieved 23 June 2015.
  8. ^ McMurray, J. (1995). Chemistry Annotated Instructors Edition (4th ed.). Prentice Hall. p. 272. ISBN 978-0-131-40221-8
  9. ^ a b c Weinhold, Frank; Landis, Clark R. (2005). Valency and bonding: A Natural Bond Orbital Donor-Acceptor Perspective. Cambridge: Cambridge University Press. pp. 367, 376, 381–383, 570–572. ISBN 978-0-521-83128-4.
  10. ^ Kaupp, Martin (2001). ""Non-VSEPR" Structures and Bonding in d(0) Systems". Angew Chem Int Ed Engl. 40 (1): 3534–3565. doi:10.1002/1521-3773(20011001)40:19<3534::AID-ANIE3534>3.0.CO;2-#.
  11. ^ Jean, Yves (2005). Molecular Orbitals of Transition Metal Complexes. Oxford: Oxford University Press. pp. 37–76. ISBN 978-0-19-151369-5.
  12. ^ a b Angelo R. Rossi; Roald. Hoffmann (1975). "Transition metal pentacoordination". Inorganic Chemistry. 14 (2): 365–374. doi:10.1021/ic50144a032.
  13. ^ a b Roald. Hoffmann; Barbara F. Beier; Earl L. Muetterties; Angelo R. Rossi (1977). "Seven-coordination. A molecular orbital exploration of structure, stereochemistry, and reaction dynamics". Inorganic Chemistry. 16 (3): 511–522. doi:10.1021/ic50169a002.
  14. ^ E. Magnusson. Hypercoordinate molecules of second-row elements: d functions or d orbitals? J. Am. Chem. Soc. 1990, 112, 7940–7951. doi:10.1021/ja00178a014
  15. ^ David L. Cooper; Terry P. Cunningham; Joseph Gerratt; Peter B. Karadakov; Mario Raimondi (1994). "Chemical Bonding to Hypercoordinate Second-Row Atoms: d Orbital Participation versus Democracy". Journal of the American Chemical Society. 116 (10): 4414–4426. doi:10.1021/ja00089a033.
  16. ^ Frenking, Gernot; Shaik, Sason, eds. (May 2014). "Chapter 7: Chemical bonding in Transition Metal Compounds". The Chemical Bond: Chemical Bonding Across the Periodic Table. Wiley -VCH. ISBN 978-3-527-33315-8.
  17. ^ C. R. Landis; F. Weinhold (2007). "Valence and extra-valence orbitals in main group and transition metal bonding". Journal of Computational Chemistry. 28 (1): 198–203. doi:10.1002/jcc.20492. PMID 17063478.
  18. ^ a b Richard D. Harcourt; Thomas M. Klapötke (2003). "Increased valence (qualitative valence bond) descriptions of the electronic structures of electron-rich fluorine-containing molecules". Journal of Fluorine Chemistry. 123 (1): 5–20. doi:10.1016/S0022-1139(03)00012-5.
  19. ^ Frenking, Gernot; Shaik, Sason, eds. (2014). "Chapter 3: The NBO View of Chemical Bonding". The Chemical Bond: Fundamental Aspects of Chemical Bonding. John Wiley & Sons. ISBN 978-3-527-66471-9.
  20. ^ Kaupp, Martin (2007). "The role of radial nodes of atomic orbitals for chemical bonding and the periodic table". Journal of Computational Chemistry. 28 (1): 320–325. doi:10.1002/jcc.20522. ISSN 0192-8651. PMID 17143872.
  21. ^ Kaupp, Martin (2014) [1st. Pub. 2014]. "Chapter 1: Chemical bonding of main group elements". In Frenking, Gernod & Shaik, Sason (eds.). The Chemical Bond: Chemical Bonding Across the Periodic Table. Wiley-VCH. ISBN 978-1-234-56789-7.
  22. ^ Kutzelnigg, W. (August 1988). "Orthogonal and non-orthogonal hybrids". Journal of Molecular Structure: THEOCHEM. 169: 403–419. doi:10.1016/0166-1280(88)80273-2.
  23. ^ Sason S. Shaik; Phillipe C. Hiberty (2008). A Chemist's Guide to Valence Bond Theory. New Jersey: Wiley-Interscience. pp. 104–106. ISBN 978-0-470-03735-5.
  24. ^ Weinhold, Frank; Landis, Clark R. (2012). Discovering Chemistry with Natural Bond Orbitals. Hoboken, N.J.: Wiley. pp. 67–68. ISBN 978-1-118-11996-9.

External links

Aluminium(I) oxide

Aluminium(I) oxide is a compound of aluminium and oxygen with the chemical formula Al2O. It can be prepared by heating the stable oxide Al2O3 with elemental silicon at 1800 °C under vacuum.

Bent's rule

In chemistry, Bent's rule describes and explains the relationship between the orbital hybridization of central atoms in molecules and the electronegativities of substituents. The rule was stated by Henry Bent as follows:

Atomic s character concentrates in orbitals directed toward electropositive substituents.

The chemical structure of a molecule is intimately related to its properties and reactivity. Valence bond theory proposes that molecular structures are due to covalent bonds between the atoms and that each bond consists of two overlapping and typically hybridised atomic orbitals. Traditionally, p-block elements in molecules are assumed to hybridise strictly as spn, where n is either 1, 2, or 3. In addition, the hybrid orbitals are all assumed to be equivalent (i.e. the n + 1 spn orbitals have the same p character). Results from this approach are usually good, but they can be improved upon by allowing isovalent hybridization, in which the hybridised orbitals may have noninteger and unequal p character. Bent's rule provides a qualitative estimate as to how these hybridised orbitals should be constructed. Bent's rule is that in a molecule, a central atom bonded to multiple groups will hybridise so that orbitals with more s character are directed towards electropositive groups, while orbitals with more p character will be directed towards groups that are more electronegative. By removing the assumption that all hybrid orbitals are equivalent spn orbitals, better predictions and explanations of properties such as molecular geometry and bond strength can be obtained. Bent's rule has been proposed as an alternative to VSEPR theory as an elementary explanation for observed molecular geometries of simple molecules with the advantages of being more easily reconcilable with modern theories of bonding and having stronger experimental support.

Bent's rule can be generalized to d-block elements as well. The hybridisation of a metal center is arranged so that orbitals with more s character are directed towards ligands that form bonds with more covalent character. Equivalently, orbitals with more d character are directed towards groups that form bonds of greater ionic character. The validity of Bent's rule for 75 bond types between the main group elements was examined recently. For bonds with the larger atoms from the lower periods, trends in orbital hybridization depend strongly on both electronegativity and orbital size.

Bent molecular geometry

In chemistry, the term "bent" can be applied to certain molecules to describe their molecular geometry. Certain atoms, such as oxygen, will almost always set their two (or more) covalent bonds in non-collinear directions due to their electron configuration. Water (H2O) is an example of a bent molecule, as well as its analogues. The bond angle between the two hydrogen atoms is approximately 104.45°. Nonlinear geometry is commonly observed for other triatomic molecules and ions containing only main group elements, prominent examples being nitrogen dioxide (NO2), sulfur dichloride (SCl2), and methylene (CH2).

This geometry is almost always consistent with VSEPR theory, which usually explains non-collinearity of atoms with a presence of lone pairs. There are several variants of bending, where the most common is AX2E2 where two covalent bonds and two lone pairs of the central atom (A) form a complete 8-electron shell. They have central angles from 104° to 109.5°, where the latter is consistent with a simplistic theory which predicts the tetrahedral symmetry of four sp3 hybridised orbitals. The most common actual angles are 105°, 107°, and 109°: they vary because of the different properties of the peripheral atoms (X).

Other cases also experience orbital hybridisation, but in different degrees. AX2E1 molecules, such as SnCl2, have only one lone pair and the central angle about 120° (the centre and two vertices of an equilateral triangle). They have three sp2 orbitals. There exist also sd-hybridised AX2 compounds of transition metals without lone pairs: they have the central angle about 90° and are also classified as bent.

Beryllium oxide

Beryllium oxide (BeO), also known as beryllia, is an inorganic compound with the formula BeO. This colourless solid is a notable electrical insulator with a higher thermal conductivity than any other non-metal except diamond, and exceeds that of most metals. As an amorphous solid, beryllium oxide is white. Its high melting point leads to its use as a refractory material. It occurs in nature as the mineral bromellite. Historically and in materials science, beryllium oxide was called glucina or glucinium oxide.

Double bond

A double bond in chemistry is a chemical bond between two chemical elements involving four bonding electrons instead of the usual two. The most common double bond occurs between two carbon atoms and can be found in alkenes. Many types of double bonds exist between two different elements. For example, in a carbonyl group with a carbon atom and an oxygen atom. Other common double bonds are found in azo compounds (N=N), imines (C=N) and sulfoxides (S=O). In skeletal formula the double bond is drawn as two parallel lines (=) between the two connected atoms; typographically, the equals sign is used for this. Double bonds were first introduced in chemical notation by Russian chemist Alexander Butlerov.Double bonds involving carbon are stronger than single bonds and are also shorter. The bond order is two. Double bonds are also electron-rich, which makes them potentially more reactive in the presence of a strong electron acceptor (as in addition reactions of the halogens).

Glossary of chemistry terms

This glossary of chemistry terms is a list of terms and definitions relevant to chemistry, including chemical laws, diagrams and formulae, laboratory tools, glassware, and equipment. Chemistry is a physical science concerned with the composition, structure, and properties of matter, as well as the changes it undergoes during chemical reactions; it has an extensive vocabulary and a significant amount of jargon.

Note: All periodic table references refer to the IUPAC Style of the Periodic Table.

Hybridisation

Hybridisation (or hybridization) may refer to:

Hybridisation (biology), the process of combining different varieties of organisms to create a hybrid

Orbital hybridisation, in chemistry, the mixing of atomic orbitals into new hybrid orbitals

Nucleic acid hybridization, the process of joining two complementary strands of nucleic acids - RNA, DNA or oligonucleotides

In evolutionary algorithms, the merging two or more optimization techniques into a single algorithm

Memetic algorithm, a common template for hybridisation

In linguistics, the process of one variety blending with another variety

The alteration of a vehicle into a hybrid electric vehicle

In globalization theory, the ongoing blending of cultures

Hybridization in political election campaign communication, the combining of campaign techniques developed in different countries

In paleoanthropology, the controversial hypothesis of Neanderthal and human hybridization

Linear acetylenic carbon

Linear acetylenic carbon (LAC), also called carbyne, is an allotrope of carbon that has the chemical structure (−C≡C−)n as a repeating chain, with alternating single and triple bonds. It would thus be the ultimate member of the polyyne family.

This polymeric carbyne is of considerable interest to nanotechnology as its Young's modulus is 32.7 TPa – forty times that of diamond. It has also been identified in interstellar space; however, its existence in condensed phases has been contested recently, as such chains would crosslink exothermically (and perhaps explosively) if they approached each other.

Linus Pauling

Linus Carl Pauling (; February 28, 1901 – August 19, 1994) was an American chemist, biochemist, peace activist, author, educator, and husband of American human rights activist Ava Helen Pauling. 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 and molecular biology. His contributions to the theory of the chemical bond include the concept of orbital hybridisation and the first accurate scale of electronegativities of the elements. Pauling also worked on the structures of biological molecules, and showed the importance of the alpha helix and beta sheet in protein secondary structure. Pauling's approach combined methods and results from X-ray crystallography, molecular model building and quantum chemistry. His discoveries inspired the work of James Watson, Francis Crick, and Rosalind Franklin on the structure of DNA, which in turn made it possible for geneticists to crack the DNA code of all organisms.In his later years he promoted nuclear disarmament, as well as orthomolecular medicine, megavitamin therapy, and dietary supplements. None of the latter have gained much acceptance in the mainstream scientific community.For his scientific work, Pauling was awarded the Nobel Prize in Chemistry in 1954. For his peace activism, he was awarded the Nobel Peace Prize in 1962. He is one of four individuals to have won more than one Nobel Prize (the others being Marie Curie, John Bardeen and Frederick Sanger). Of these, he is the only person to have been awarded two unshared Nobel Prizes, and one of two people to be awarded Nobel Prizes in different fields, the other being Marie Curie.

Molecular geometry

Molecular geometry is the three-dimensional arrangement of the atoms that constitute a molecule. It includes the general shape of the molecule as well as bond lengths, bond angles, torsional angles and any other geometrical parameters that determine the position of each atom.

Molecular geometry influences several properties of a substance including its reactivity, polarity, phase of matter, color, magnetism and biological activity. The angles between bonds that an atom forms depend only weakly on the rest of molecule, i.e. they can be understood as approximately local and hence transferable properties.

Nicholas reaction

The Nicholas reaction is an organic reaction where a dicobalt octacarbonyl-stabilized propargylic cation is reacted with a nucleophile. Oxidative demetallation gives the desired alkylated alkyne. It is named after Kenneth M. Nicholas.

Several reviews have been published.

Nitrogen

Nitrogen is the chemical element with the symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. The name nitrogène was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790, when it was found that nitrogen was present in nitric acid and nitrates. Antoine Lavoisier suggested instead the name azote, from the Greek ἀζωτικός "no life", as it is an asphyxiant gas; this name is instead used in many languages, such as French, Russian, Romanian and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.

Nitrogen is the lightest member of group 15 of the periodic table, often called the pnictogens. The name comes from the Greek πνίγειν "to choke", directly referencing nitrogen's asphyxiating properties. It is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N2. Dinitrogen forms about 78% of Earth's atmosphere, making it the most abundant uncombined element. Nitrogen occurs in all organisms, primarily in amino acids (and thus proteins), in the nucleic acids (DNA and RNA) and in the energy transfer molecule adenosine triphosphate. The human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. The nitrogen cycle describes movement of the element from the air, into the biosphere and organic compounds, then back into the atmosphere.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong triple bond in elemental nitrogen (N≡N), the second strongest bond in any diatomic molecule after carbon monoxide (CO), dominates nitrogen chemistry. This causes difficulty for both organisms and industry in converting N2 into useful compounds, but at the same time means that burning, exploding, or decomposing nitrogen compounds to form nitrogen gas releases large amounts of often useful energy. Synthetically produced ammonia and nitrates are key industrial fertilisers, and fertiliser nitrates are key pollutants in the eutrophication of water systems.

Apart from its use in fertilisers and energy-stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric and cyanoacrylate used in superglue. Nitrogen is a constituent of every major pharmacological drug class, including antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by metabolizing into nitric oxide. Many notable nitrogen-containing drugs, such as the natural caffeine and morphine or the synthetic amphetamines, act on receptors of animal neurotransmitters.

SP3

SP3 may refer to:

sp3 hybrids, a type of orbital hybridisation

Sp3 transcription factor, a protein and the gene it is encoded

Savoia-Pomilio SP.3, a reconnaissance and bomber aircraft built in Italy

USS Zipalong (SP-3), an armed motorboat

1971 SP3 or 3922 Heather, a main-belt asteroid discovered on September 26, 1971

1984 SP3 or 3155 Lee, a main-belt asteroid discovered on September 28, 1984

SP3, a model of steam toy made by British manufacturer Mamod

Service pack 3, for computer software

Socket SP3, a CPU socket for AMD processors

SP3, a sink in the Sima Pumacocha, a cave in Peru

Surface Pro 3, a 2-in-1 personal computer by Microsoft

Spinterface

Spinterface is a term coined to indicate an interface between a ferromagnet and an organic semiconductor.

This is a widely investigated topic in molecular spintronics, since the role of interfaces plays a huge part in the functioning of a device. In particular, spinterfaces are widely studied in the scientific community because of their hybrid organic/inorganic composition. In fact, the hybridization between the metal and the organic material can be controlled by acting on the molecules, which are way more responsive to electrical and optical stimuli with respect to metals. This gives rise to the possibility of efficiently tuning the magnetic properties of the interface at the atomic scale.

Tetrahedral molecular geometry

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(−⅓) = 109.4712206...° ≈ 109.5° when all four substituents are the same, as in methane (CH4) as well as its heavier analogues. Methane and other perfectly symmetrical tetrahedral molecules belong to point group Td, but most tetrahedral molecules have lower symmetry. Tetrahedral molecules can be chiral.

Titanium(IV) hydride

Titanium(IV) hydride (systematically named titanium tetrahydride) is an inorganic compound with the empirical chemical formula TiH4. It has not yet been obtained in bulk, hence its bulk properties remain unknown. However, molecular titanium(IV) hydride has been isolated in solid gas matrices. The molecular form is a colourless gas, and very unstable toward thermal decomposition. As such the compound is not well characterised, although many of its properties have been calculated via computational chemistry.

VSEPR theory

Valence shell electron pair repulsion theory, or VSEPR theory ( VESP-ər, və-SEP-ər), 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, Ronald Gillespie and Ronald Nyholm.

The premise of VSEPR is that the valence electron pairs surrounding an atom tend to repel each other and will, therefore, adopt an arrangement that minimizes this repulsion, thus determining the molecule's geometry. Gillespie has emphasized that the electron-electron repulsion due to the Pauli exclusion principle is more important in determining molecular geometry than the electrostatic repulsion.VSEPR theory is based on observable electron density rather than mathematical wave functions and hence unrelated to orbital hybridisation, although both address molecular shape. While it is mainly qualitative, VSEPR has a quantitative basis in quantum chemical topology (QCT) methods such as the electron localization function (ELF) and the quantum theory of atoms in molecules (QTAIM).

Basics
Types of bonds
Valence bond theory
Molecular orbital theory

Languages

This page is based on a Wikipedia article written by authors (here).
Text is available under the CC BY-SA 3.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.