Chemical bond
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A chemical bond is the physical phenomenon (or phenomena) responsible for the attractive interactions between atoms that confers stability to di- and polyatomic chemical compounds. The explanation of the attractive forces is a complex area that is described by the laws of quantum electrodynamics, but usually chemists rely on quantum theory or qualitative descriptions that are less rigorous but more easily explained. In general, strong chemical bonding is associated with the sharing or transfer of electrons between the participating atoms. Molecules, crystals, and diatomic gases - indeed most of the physical environment around us - are held together by chemical bonds, which dictate the structure of matter.
Bonds vary widely in their strength. Weak interactions between atoms and molecules can arise from induced polarity (such as London forces) between the electron clouds. Generally covalent and ionic bonds are often described as strong, whereas hydrogen bonds and van der Waals are generally considered to be weaker. Care should be taken because the strongest of the weak cases can be stronger than the weakest of the "strong" bonds.
[edit] Overview
In the simplest view of a so-called covalent bond, one or more electrons - often a pair - is drawn into the space between two atomic nuclei, because in this region the negatively charged electrons are attracted to the positive charge of both nuclei, instead of just one of them. This overcomes the repulsion between nuclei. This situation holds the nuclei in a relatively fixed equilibrium configuration, although they will still vibrate about that equilibrium position. In summary, covalent bonding involves sharing of electrons in which the positively charged nuclei of two or more atoms simultaneously attract the negatively charged electrons that are being shared.
In a simplified view of an ionic bond, one or more electrons is simply transferred from one atom to another, causing one atom to assume net positive charge, and the other a negative charge. The bond then results from electrostatic attraction between atoms.
All bonds can be explained by quantum theory, but in practice, simplification rules allow chemists to predict the strength, directionality, and polarity of bonds. The octet rule and VSEPR theory are two examples. More sophisticated theories are valence bond theory which includes orbital hybridization and resonance, and the linear combination of atomic orbitals molecular orbital method which includes ligand field theory. Electrostatics are used to describe bond polarities and the effects they have on chemical substances.
[edit] History
Early speculations into the nature of the chemical bond, from as early as the 12th century, supposed that certain types of chemical species were joined by a type of chemical affinity. By the mid 19th century, Edward Frankland, F.A. Kekule, A.S. Couper, A.M. Butlerov, and Hermann Kolbe, building on the theory of radicals, developed the theory of valency, originally called “combining power”, in which compounds were joined owing to an attraction of positive and negative poles. In 1916, chemist Gilbert Lewis developed the concept of the electron-pair bond.
In 1927, the first mathematically complete quantum description of a simple chemical bond, i.e. that produced by one electron in the hydrogen molecular ion, H2+, was derived by the Danish physicist Oyvind Burrau.[1] This work showed that the quantum approach to chemical bonds could be fundamentally and quantitatively correct, but the mathematical methods used could not be extended to molecules containing more than one electron. A more practical, albeit less quantitative, approach was put forward in the same year by Walter Heitler and Fritz London. The Heitler-London method forms the basis of what is now called valence bond theory. In 1929, the linear combination of atomic orbitals molecular orbital method (LCAO) approximation was introduced by Sir John Lennard-Jones, who also suggested methods to derive electronic structures of molecules of F2 (fluorine) and O2 (oxygen) molecules, from basic quantum principles. This molecular orbital theory represented a covalent bond as a orbitals formed by combining the quantum mechanical Schrödinger atomic orbitals which had been hypothesized for electrons in single atoms. The equations for bonding electrons in multi-electron atoms could not be solved to mathematical perfection (i.e., analytically), but approximations for them still gave many good qualitative preditions and results. Most quantitative calculations in modern quantum chemistry use either valence bond or molecular orbital theory as a starting point, although a third approach, Density Functional Theory, has become increasingly popular in recent years.
In 1935, H. H. James and A. S. Coolidge carried out a calculation on the dihydrogen molecule that, unlike all previous calculation which used functions only of the distance of the electron from the atomic nucleus, used functions which also explicitly added the distance between the two electrons.[2] With up to 13 adjustable parameters they obtained a result very close to the experimental result for the dissociation energy. Later extensions have used up to 54 parameters and give excellent agreement with experiment. This calculation convinced the scientific community that quantum theory could give agreement with experiment. However this approach has none of the physical pictures of the valence bond and molecular orbital theories and is difficult to extend to larger molecules.
[edit] Valence bond theory
In 1931, chemist Linus Pauling published what some consider one of the most important papers in the history of chemistry: “On the Nature of the Chemical Bond”. In this paper, building on the works of Lewis, and the valence bond theory (VB) of Heitler and London, and his own earlier work, he presented six rules for the shared electron bond, the first three of which were already generally known:
- [1] the electron-pair bond forms through the interaction of an unpaired electron on each of two atoms.
- [2] the spins of the electrons have to be opposed.
- [3] once paired, the two electrons cannot take part in additional bonds.
His last three rules were new:
- [4] the electron-exchange terms for the bond involves only one wave function from each atom.
- [5] the available electrons in the lowest energy level form the strongest bonds.
- [6] of two orbitals in an atom, the one that can overlap the most with an orbital from another atom will form the strongest bond, and this bond will tend to lie in the direction of the concentrated orbital.
Building on this article, Pauling’s 1939 textbook: On the Nature of the Chemical Bond would become what some have called the “bible” of modern chemistry. This book helped experimental chemists to understand the impact of quantum theory on chemistry. However, the later edition in 1959 failed to address adequately the problems that appeared to be better understood by molecular orbital theory. The impact of valence theory declined during the 1960's and 1970's as molecular orbital theory grew in popularity and was implemented in many large computer programs. Since the 1980s, the more difficult problems of implementing valence bond theory into computer programs have been largely solved and valence bond theory has seen a resurgance.
[edit] Molecular orbital theory
Molecular orbital theory (MO) uses a linear combination of atomic orbitals to form molecular orbitals which cover the whole molecule. These are often divided into bonding orbitals, anti-bonding orbitals, and non-bonding orbitals. A molecular orbital is merely a Schrödinger orbital which includes several, but often only two nuclei. If this orbital is of type in which the electron(s) in the orbital have a higher probability of being between nuclei than elsewhere, the orbital will be a bonding orbital, and will tend to hold the nuclei together. If the electrons tend to be present in a molecular orbital in which they spend more time elsewhere than between the nuclei, the orbital will function as an anti-bonding orbital and will actually weaken the bond. Electrons in non-bonding orbitals tend to be in deep orbitals (nearly atomic orbitals) associated almost entirely with one nucleus or the other, and thus they spend equal time between nuclei or not. These electrons neither contribute nor detract from bond strength.
Molecular orbitals are further divided according the types of atomic orbitals combining to form a bond. These orbitals are results of electron-nucleus interactions that are caused by the fundamental force of electromagnetism. Chemical substances will form a bond if their orbitals become lower in energy when they interact with each other. Different chemical bonds are distinguished that differ by electron cloud shape and by energy levels.
MO theory provides a global, delocalized perspective on chemical bonding. For example, in the MO theory for hypervalent molecules, it is no longer necessary to invoke a major role for d-orbitals. In MO theory, any electron in a molecule may be found anywhere in the molecule, since quantum conditions allow electrons to travel under the influence of an arbitrarily large number of nuclei, so long as permitted by certain quantum rules. Athough in MO theory some molecular orbitals may hold electrons which are more localized between specific pairs of molecular atoms, other orbitals may hold electrons which are spread more uniformly over the molecule. Thus, overall, bonding (and electrons) are far more delocalized (spread out) in MO theory, than is implied in VB theory. This makes MO theory more useful for the description of extended systems.
An example is that in the MO picture of benzene, composed of a hexagonal ring of 6 carbon atoms. In this molecule, 24 of the 30 total valence bonding electrons are located in 12 σ (sigma) bonding orbitals which are mostly located between pairs of atoms (C-C or C-H), similar to the valence bond picture. However, in benzene the remaining 6 bonding electrons are located in 3 π (pi) molecular bonding orbitals that are delocalized around the ring. Two are in a MO which has equal contrinutions from all 6 atoms. The other two have a vertical nodes at right angles to each other. As in the VB theory, all of these 6 delocalized pi electrons reside in a larger space which exists above and below the ring plane. All carbon-carbon bonds in benzene are chemically equivalent. In MO theory this is a direct consequence of the fact that the 3 molecular pi orbitals form a combination which evenly spreads the extra 6 electrons over 6 carbon atoms. [1]
In molecules such as methane, the 8 valence electrons are in 4 MOs that are spread out over all 5 atoms. However, it is possible to transform this picture, without altering the total wavefunction and energy, to one with 8 electrons in 4 localised orbitals that are similar to the normal bonding picture of four two-electron covalent bonds. This is what has been done above for the σ (sigma) bonds of benzene, but it is not possible for the π (pi) orbitals. The delocalised picture is more appropriate for ionisation and spectroscopic properties. Upon ionization, a single electron is taken from the whole molecule. The resulting ion does not have one bond different from the other three Similarly for electronic excitations, the electron that is excited is found over the whole molecule and not in one bond.
As in benzene, in substances such as beta carotene, chlorophyll or heme, some electrons the π (pi) orbitals are spread out in molecular orbitals over long distances in a molecule, giving rise to light absorption in lower energies (visible colors), a fact which is observed. This and other spectroscopic data for molecules are better explained in MO theory, with an emphasis on electronic states associated with multicenter orbitals, including mixing of orbitals premised on principles of orbital symmetry matching. The same MO principles also more naturally explain some electrical phenomena, such as high electrical conductivity in the planar direction of the hexagonal atomic sheets that exist in graphite. In MO theory, "resonance" (a mixing and blending of VB bond states) is a natural consequence of symmetry. For example, in graphite, as in benzene, it is not necessary to invoke the sp2 hybridization and resonance of VB theory, in order to explain electrical conduction. Instead, MO theory simply recognizes that some electrons in the graphic atomic sheets are completely delocalized over arbitrary distances, and reside in very large molecular orbitals that cover an entire graphite sheet, and some electrons are thus are as free to move and conduct electricity in the sheet plane, as if they resided in a metal.
[edit] Comparision of valence bond and molecular orbital theory
In some respects valence bond theory is superior to molecular orbital theory. When applied to the simplest two-electron molecule, H2, valence bond theory, even at the simplest Heitler-London approach, gives a much closer approximation to the bond energy, and it provides a much more accurate representation of the behavior of the electrons as chemical bonds are formed and broken. In contrast simple molecular orbital theory predicts that the hydrogen molecule dissociates into a linear superposition of hydrogen atoms and positive and negative hydrogen ions, a completely unphysical result. This explains in part why the curve of total energy against interatomic distance for the valence bond method lies above the curve for the molecular orbital method at all distances and most particularly so for large distances. This situation arises for all homonuclear diatomic molecules and is particularly a problem for F2, where the minimum energy of the curve with molecular orbital theory is still higher in energy than the energy of two F atoms.
The concepts of hybridization are so versatile, and the variability in bonding in most organic compounds is so modest, that valence bond theory remains an integral part of the vocabulary of organic chemistry. However, the work of Friedrich Hund, Robert Mulliken, and Gerhard Herzberg showed that molecular orbital theory provided a more appropriate description of the spectroscopic, ionization and magnetic properties of molecules. The deficiencies of valence bond theory became apparent when hypervalent molecules (e.g. PF5) were explained without the use of d orbitals that were were crucial to the bonding hybidisation scheme proposed for such molecules by Pauling. Metal complexes and electron deficient compounds (e.g. diborane) also appeared to be well described by molecular orbital theory, although valence bond descriptions have been made.
In the 1930s the two methods strongly competed until it was realised that they are both approximations to a better theory. If we take the simple valence bond structure and mix in all possible covalent and ionic structures arising from a particular set of atomic orbitals, we reach what is called the full configuration interaction wave function. If we take the simple molecular orbital description of the ground state and combine that function with the functions describing all possible excited states using unoccupied orbitals arising from the same set of atomic orbitals, we also reach the full configuration interaction wavefunction. It can be then seen that the simple molecular orbital approach gives too much weight to the ionic structures, while the simple valence bond approach gives too little. This can also be described as saying that the molecular orbital approach is too delocalised, while the valence bond approach is too localised.
The two approaches are now regarded as complementary, each providing its own insights into the problem of chemical bonding. Modern calculations in quantum chemistry usually start from (but ultimately go far beyond) a molecular orbital rather than a valence bond approach, not because of any intrinsic superiority in the former but rather because the MO approach is more readily adapted to numerical computations. However better valence bond programs are now available.
[edit] Bonds in chemical formulas
The 3-dimensionality of atoms and molecules makes it difficult to use a single technique for indicating orbitals and bonds. In molecular formulas the chemical bonds (binding orbitals) between atoms are indicated by various different methods according to the type of discussion. Sometimes, they are completely neglected. For example, in organic chemistry chemists are sometimes concerned only with the functional groups of the molecule. Thus, the molecular formula of ethanol (a compound in alcoholic beverages) may be written in a paper in conformational, 3-dimensional, full 2-dimensional (indicating every bond with no 3-dimensional directions), compressed 2-dimensional (CH3-CH2-OH), separating the functional group from another part of the molecule (C2H5OH), or by its atomic constituents (C2H6O), according to what is discussed. Sometimes, even the non-bonding valence shell electrons ( with the 2-dimensionalized approximate directions) are marked, f.e. for elemental carbon .'C.' Some chemists may also mark the respective orbitals, f.e. the hypothetical ethene-4 anion (\/C=C/\ -4) indicating the possibility of bond formation.
[edit] Strong chemical bonds
These chemical bonds are intramolecular forces, which hold atoms together in molecules and in solids. In the simplistic localized view of bonding, the number of electrons participating in a bond (or located in a bonding orbital) is typically multiples of two, four, or six, respectively. Even numbers are common because electrons enjoy lower energy states, if paired. Substantially more advanced bonding theories have shown that bond strength is not always a whole number, depending on the distribution of electrons to each atom involved in a bond. For example, the carbons in benzene are connected to each other with about 1.5 bonds, and the two atoms in nitric oxide NO, are connected with about 2.5 bonds. Quadruple bonds are also well known. The type of strong bond depends on the difference in electronegativity and the distribution of the electron orbital paths available to the atoms that are bonded. The larger the electronegativity, the more an electron is attracted to a particular atom involved in the bond, and the more "ionic" properties the bond is said to have ("ionic" means the bond electron(s) are unequally shared). The smaller the electronegativity, the more covalent properties (full sharing) the bond has.
[edit] Covalent bond
Covalent bonding is a common type of bonding, in which the electronegativity difference between the bonded atoms is small or non-existent. Bonds within most organic compounds are described as covalent. See sigma bonds and pi bonds for LCAO-description of such bonding.
[edit] Polar covalent bond
Polar covalent bonding is intermediate in character between a covalent and an ionic bond.
[edit] Ionic bond
Ionic bonding is a type of electrostatic interaction between atoms which have an electronegativity difference of over 1.6 (this limit is a convention).[citation needed] These form in a solution between two ions after the excess of the solvent is removed. Ionic charges are commonly between -3e to +7e
[edit] Other strong bonds
[edit] Coordinate covalent bond
Coordinate covalent bonding is a kind of covalent bonding, in which the covalent bonding electrons originate solely from one of the atoms, the electron-pair donor or Lews base but are approximately equally shared in the formation of a covalent bond. This concept is somewhat fading as chemists increasingly embrace molecular orbital theory. Examples of coordinate covalent bonding occur in nitrones and ammonia-borane. The arrangement is different from an ionic bond in that the electronegativity difference is small, resulting in covalency.
[edit] Banana bond
The Banana bond is a kind of bonding in which the bond bends, often due to the presence of an influencing atom in the middle of another covalent bond. These bonds are likely to be more susceptible to reactions than ordinary bonds. An example can be found in diborane, where the bonds between boron atoms are known as "three center, two-electron bonds". Each such bond (2 per molecule in diborane) contains a pair of electrons which connect the boron atoms to each other in a banana shape, with a proton (nucleus of a hydrogen atom) in the middle of the bond, sharing electrons with both boron atoms.
[edit] Chemical bonds involving more than two atoms
[edit] Aromatic bond
In most cases, the locations of electrons cannot be simplified to simple lines (place for two electrons) or dots (a single electron). In aromatic bonds which occur in rings of atoms where the 4n+2 rule determines whether ring molecules comprised of C=C bonds would show behavior extra stability by allowing extra sharing of electrons below and above the ring plane.
In benzene, the prototypical aromatic compound, 18 bonding electrons bind 6 carbon atoms together to form a planar ring structure. The bond "order" (average number of bonds) between the different carbon atoms may be said to be (18/6)/2=1.5, but in this case the bonds are all identical from the chemical point of view. They may sometimes be written as single bonds alternating with double bonds, but the view of all ring bonds as being equivalently about 1.5 bonds in strength, is much closer to truth.
In the case of heterocyclic aromatics and substituted benzenes, the electronegativity differences between different parts of the ring may dominate the chemical behaviour of aromatic ring bonds, which otherwise are equivalent.
[edit] Metallic bond
In a metallic bond, bonding electrons are delocalized over a lattice of atoms. By contrast, in ionic compounds, the locations of the binding electrons and their charges are static.
[edit] Intermolecular bonding
There are four basic types of bonds that can be formed between two or more (otherwise non-associated) molecules, ions or atoms.
[edit] Permanent dipole to permanent dipole
A large electronegativity difference between two strongly bonded atoms within a molecule causes a dipole to form (a dipole is a pair of permanent partial charges). Dipoles will attract or repel each other.
[edit] Hydrogen bond
In some ways this is an especially strong example of a permanent dipole bond, as above. However, in the hydrogen bond, the hydrogen proton comes closer to being shared between target and donor atoms, in a three-center system like the banana bond in diborane. Hydrogen bonds explain the relatively high boiling points of liquids like water, ammonia, and hydrogen fluoride, compared with their heavier counterparts in the same periodic table column.
[edit] Instantaneous dipole to induced dipole (van der Waals)
Instantaneous dipole to induced dipole, or van der Waals forces, are the weakest, but also the most prolific - occurring between all chemical substances. Imagine a helium atom: At any one point in time, the electron cloud around the - otherwise-neutral - atom can be thought to be slightly imbalanced, with momentarily more negative charge on one side. This is referred to as an instantaneous dipole. This dipole, with its slight charge imbalance, may attract or repel the electrons within a neighbouring helium atom, setting up another dipole. The two atoms will be attracted for an instant, before the charge rebalances and the atoms move on.
[edit] Cation-pi interaction
Cation-pi interactions occur between the localized negative charge of π orbital electrons, located above and below the plane of an aromatic ring, and a positive charge.
[edit] Electrons in chemical bonds
Many simple compounds involve covalent bonds. These molecules have structures that can be predicted using valence bond theory, and the properties of atoms involved can be understood using concepts such as oxidation number. Other compounds that involve ionic structures can be understood using theories from classical physics.
In the case of ionic bonding, electrons are mainly localized on the individual atoms, and electrons do not travel between the atoms very much. Each atom is assigned an overall electric charge to help conceptualize the molecular orbital's distribution. The forces between atoms (or ions) are largely characterized by isotropic continuum electrostatic potentials.
By contrast, in covalent bonding, the electron density within a bond is not assigned to individual atoms, but is instead delocalized in the MOs between atoms. The widely-accepted theory of the linear combination of atomic orbitals (LCAO) helps describe the molecular orbital structures and energies based on the atomic orbitals of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties. These may have their own names, too, such as Sigma and Pi bond.
Atoms can also form bonds that are intermediates between ionic and covalent. This is because these definitions are based on the extent of electron delocalization. Electrons can be partially delocalized between atoms, but spend more time around one atom than another. This type of bond is often called polar covalent. See electronegativity.
Thus, the electrons in a molecular orbital (or 'in a polar covalent, or in a covalent bond') can be said to be either localized on certain atom(s) or delocalized between two or more atoms. The type of bond between two atoms is defined by how much the electron density is localized or delocalized among the atoms of the substance.
[edit] Determination of chemical properties through chemical bonding
Intermolecular forces cause molecules to be attracted or repulsed by each other. Often, these define some of the physical characteristics (such as the melting point) of a substance. These forces include ionic interactions, hydrogen bonds, dipole-dipole interactions, and induced dipole interactions.
[edit] See also
- Electron
- Electronegativity
- Periodic table
- octet rule
- delocalized electron
- Valence shell
- Ion
- Valence bond theory
- hybridization
- sigma, pi and delta bonds
- Chemical reaction
More advanced articles:
- Bohr model
- quantum number
- List of Hund's rules
- Quantum chemistry
- LCAO
- Atomic orbital, molecular orbital
[edit] References
- ^ Laidler, K. J. (1993) The World of Physical Chemistry, Oxford University Press, p. 347
- ^ (Journal of Chemical Physics, 1, pg 823, 1933)
- W. Locke (1997). Introduction to Molecular Orbital Theory. Retrieved May 18, 2005.
- Carl R. Nave (2005). HyperPhysics. Retrieved May 18, 2005.
