Chemical bond

A chemical bond is the phenomenon of atoms being held together to form molecules or crystals. All chemical bonds result from electrostatic interactions and/or from sharing of electrons, and derive from the action of the electromagnetic force. Where the electromagnetic force is one of the four fundamental interactions of the physical universe. Electrostatic forces arise from the coulombic attraction or repulson between charged particles. On the other hand the sharing of electrons is described by quantum mechanics in Valence bond theory and Molecular orbital theory.

There are 5 different classes of chemical bonds that we use to categorize bonding. Actual bonds may have have properties that aren't so discretely categorized, so a given bond could be defined by more than one of these terms.

The Ionic Bond

Original article: ionic bond

The Sodium Chloride Crystal Structure is an example of a crystal lattice held together by electrostatic forces. Each atom has six nearest neighbors, with octahedral geometry. This arrangement is known as cubic close packed (ccp).
Light blue = Na⁺ ion
Dark green = Cl⁻ ion

The ionic bond refers to the eletrostatic forces that hold the ions together in an ionic solid. Ionic solids are electrically neutral and contain both cations (positively charged ions), and anions (negatively charged ions). These ions are held together in an ordered three dimensional array, a crystal lattice, by the coulombic attractive and repulsive forces acting between charges on the ions. The lattice minimizes the repulsive forces between the similarly charged ions and maximizes the attractive forces between the oppositely charged ions. While ions can be monatomic or polyatomic one way to think about the formation of an ionic bond is to look at the reaction between a metal and a non-metal; such as that between lithium and fluorine shown here. Metals loose electrons to form cations, and non-metals gain electrons to form anions. The resulting ionic solid will be held together by the electrostatic forces acting between the charged ions, the ionic bond.

$Li+F\ \ \ \to \ \ \ Li^{+}F^{-}\,\!$

Lithium has one valence electron. This electron is held rather loosely and is easily lost. Fluorine has seven valence electrons and it tends to gain one electron in order to give a stable octet. If the electron moves from lithium to fluorine each ion acquires the electronic configuration of a noble gas. The bonding energy (from the electrostatic forces acting within the crystal structure) is large enough that the overall bonded state energy is lower than the unbonded state. These type of reactions only occurs if the overall energy change for the reaction is favourable (the bonded atoms have a lower energy than the free ones). The larger the energy change the stronger the bond.

The electrostatic forces acting within an ionic solid are quite large. Such bonds are stronger than hydrogen bonds, but similar in strength to covalent bonds. Consequently ionic solids have relatively high melting points and boiling points. Moreover ions are solvated well by water so many ionic compounds are water soluable. In the solid state they do not conduct electricity as the ions are fixed in place. When molten or in solution the ions are free to move and in these states ionic compounds conduct electricity via the movement of ions themselves rather than electrons.

This section was copied from the  
bond&oldid=24707153 Ionic bond article at Wikipedia.

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The Covalent Bond

Original article: covalent bond

Covalently bonded hydrogen and carbon in a molecule of methane. One way of representing covalent bonding in a molecule is with a Lewis diagram.

Covalent bonding is a form of chemical bonding characterized by the sharing of one or more pairs of electrons between two atoms, producing a mutual attraction that holds the resultant molecule or polyatomic ion together. Atoms tend to share electrons in such a way that their outer electron shells satisfy the octet rule. Such bonds are always stronger than the intermolecular hydrogen bond and similar in strength to or stronger than the ionic bond.

Covalent bonding most frequently occurs between atoms with similar electronegativities. For this reason, non-metals tend to engage in covalent bonding more readily since metals have access to metallic bonding, where the easily-removed electrons are freer to roam about. For non-metals, liberating an electron is more difficult, so sharing is the only option when confronted with another species of similar electronegativity.

However, covalent bonding in metals and, particularly between metals and organic compounds is particularly important, especially in industrial catalysis and process chemistry, where many indispensible reactions depend on covalent bonding with metals.

Lewis Diagrams

Some example Lewis diagramss. Each bond, line connecting atoms, represents two shared electrons

The idea of covalent bonding can be traced to Gilbert N. Lewis, who in 1916 described the sharing of electron pairs between atoms. He introduced the so called Lewis Notation or Electron Dot Notation in which valence electrons (those in the outer shell) 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 and triple bonds. Some examples of Electron Dot Notation are shown in the figure to the left. In this diagram bond-forming electron pairs are represented as solid lines.

Bond order

Bond order is the scientific term used to describe the number of pairs of electrons shared between atoms forming a covalent bond. The most common type of covalent bond is the single bond, the sharing of only one pair of electrons between two individual atoms. All bonds with more than one shared pair are called multiple covalent bonds. The sharing of two pairs is called a double bond and the sharing of three pairs is called a triple bond. An example of a double bond is the carbon - oxygen bond in carbon dioxide (see diagram at left), and an example of a triple bond is the carbon - nitrogen bond in hydrogen cyanide.

A single bond usually consists of one sigma bond, a double bond of one sigma and one pi bond, and a triple bond of one sigma and two pi bonds.

Quadruple bonds, though rare, also exist. Both carbon and silicon can theoretically form these; however, the formed molecules are explosively unstable. Stable quadruple bonds are observed as transition metal-metal bonds, usually between two transition metal atoms in organometallic compounds. Molybdenum and Ruthenium are the elements most commonly observed with this bonding configuration. An example of a quadruple bond is also found in Di-tungsten tetra(hpp). Quintuple Bonds are found to exist in certain chromium dimers.

Sextuple bonds of order 6 have also been observed in transition metals in the gaseous phase at very low temperatures and are extremely rare.

Other more exotic bonds, such as three center bonds are known and defy the conventions of Lewis structres. It is also important to note that bond order is an integer value only in the elementary sense and is often fractional in more advanced contexts.

Resonance

Some structures can have more than one valid Lewis Dot Structure (for example, ozone, O₃). In an LDS diagram of O₃, the center atom will have a single bond with one atom and a double bond with the other. The LDS diagram cannot tell us which atom has the double bond; the first and second adjoining atoms have equal chances of having the double bond. These two possible structures are called resonance structures. In reality, the structure of ozone is a resonance hybrid between its two possible resonance structures. Instead of having one double bond and one single bond, there are actually two 1.5 bonds with approximately three electrons in each at all times.

A special resonance case is exhibited in aromatic rings of atoms (for example, benzene). Aromatic rings are composed of atoms arranged in a circle (held together by covalent bonds) that alternate between single and double bonds according to their LDS. In actuality, the electrons tend to be disambiguously and evenly spaced within the ring. Electron sharing in aromatic structures is often represented with a ring inside the circle of atoms.

Quantum Mechanics

While the idea of shared electron pairs provides an effective qualitative picture of covalent bonding, quantum mechanics is needed to understand the nature of these bonds and predict the structures and properties of simple molecules. Heitler and London are credited with the first successful quantum mechanical explanation of a chemical bond, specifically that of molecular hydrogen, in 1927. Their work was based on the valence bond model, which assumes that a chemical bond is formed when there is good overlap between the atomic orbitals of participating atoms. In valence bond theory molecular geometries are accounted for by the formation of hybrid atomic orbitals through the conbination of normal atomic orbitals. These hybrid atomic orbitals have specific angular relationships between each other, and thus the valence bond model can successfully predict the bond angles observed in simple molecules.

The valence bond model is supplemented with the molecular orbital model. In this model, as atoms are brought together, the atomic orbitals interact to form molecular orbitals that extend between the two bonding atoms. These molecular orbitals can be mathematically constructed through a linear combination of atomic orbitals in LCAO theory.

Using quantum mechanics it is possible to calculate the electronic structure, energy levels, bond angles, bond distances, dipole moments, and frequency spectra of simple molecules with a high degree of accuracy. Currently, bond distances and angles can be calculated as accurately as they can be measured (distances to a few pm and bond angles to a few degrees). For the case small molecules, energy calculations are sufficiently accurate to be useful for determining thermodynamic heats of formation and kinetic activation energy barriers.

Coordinate Covalent Bond

Original article: Coordinate covalent bond

A coordinate covalent bond (also known as dative covalent bond) is a special type of covalent bond in which the shared electrons come from one of the atoms only. Once the bond has been formed, its strength is no different from that of a covalent bond.

Coordinate covalent bonds are formed when a Lewis base (an electron donor or giver) donates a pair of electrons to a Lewis acid (an electron accepter) and the resultant compound is then called an adduct (a compound formed by the addition reaction between two molecules). The process of forming a dative bond is called coordination.

A compound that contains a lone pair of electrons is capable of forming a coordinate bond. Coordinate bonds can be found in many different substances, such as in simple molecules like carbon monoxide (CO), which contains one coordinate bond and two normal covalent bonds between the carbon atom and the oxygen atom, or the ammonium ion (NH₄⁺), where a coordinate bond is formed between a proton (a H⁺ ion) and the nitrogen atom. Coordinate bonds are also formed in electron deficient compounds, such as in solid beryllium chloride (BeCl₂), in which every beryllium atom is bonded to four chlorine atoms, two with normal covalent bonding, and the other two with coordinate bonds, which will give it a stable octet of electrons.

Coordinate bonding can also be found in coordination complexes involving metal ions, especially if they are transition metal ions. In such complexes, substances in a solution act as Lewis bases and donate their free pairs of electrons to the metal ion, which acts as a Lewis acid and accepts the electrons. Coordinate bonds form and the resulting compound is called a coordination complex, while the electron donors are called ligands. There are many chemicals with atoms that have lone pairs of electrons, often oxygen, sulfur, nitrogen and halogens or halide ions, which, in solution, can donate their electron pairs to become ligands. A common ligand is water (H₂O), which will form coordination complexes with any hydrated metal ions, like Cu²⁺, which will form [Cu(H₂O)₆]²⁺ in aqueous solution. Other common simple ligands are ammonia (NH₃), fluoride ions (F^-), chloride ions (Cl^-) and cyanide ions (CN^-).

A coordinate bond is represented by an arrow pointing from the donor of the electron pair to the acceptor of the electron pair.

External links

• Chemguide

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Metallic Bond

The traditional picture of the metallic bond was developed soon after the discovery of the electron. In this picture the valence electrons were viewed as an electron-gas permeating the crystal lattice structre of the metal atoms. The cohesive forces holding the metal together could be qualitatively seen as resulting from the electrostatic interaction of the positively charged metal atom cores with the negatively charged electron-gas.

With the development of quantum mechanics this picture has changed significantly. Quantum mechanically covalent bonding can be described by constructing molecular orbitals from a linear combination of atomic orbitals (LCAO theory). For some molecules, such as benzene, which contain resonance in an extended conjugated system, some of the molecular orbitals are delocalized and the electrons not fixed between particular atoms. Applying these concepts to metals leads to the Band theory of solids. In this theory the metallic bond is similar to the delocalized molecular orbitals of benzene, but on a much larger scale and with the bonding spread throughout the metal.

Let us consider a sodium atom. It has one valence electron in an s orbital. When we add another atom the overlap of the two s orbitals results in two molecular orbitals of differing energies. Adding a third atom results in three molecular orbitals and so on. In a small lump of sodium there are many thousands of atoms and consequently many thousands of molecular orbitals whose energies lie close together, This results in an energy band populated by the available electrons that is called the s-band. Similarly, with other atoms, you can also get a p-band from the overlap of atomic p orbitals.

In semiconductors there is a significant energy difference between different bands called a band gap. This is what gives semiconductors their usefull properties exploited in electronic circuits.

Hydrogen Bond

Original article: hydrogen bond

File:Wasserstoffbrückenbindungen des Wasser.png

Hydrogen bonds between water molecules are diagramatically represented by the black lines. The red lines are covalent bonds that hold oxygen (red) and hydrogen (blue) atoms together in the water molecules.

In chemistry, a hydrogen bond is a type of attractive intermolecular force that exists between two partial electric charges of opposite polarity. Although stronger than most other intermolecular forces, the typical hydrogen bond is much weaker than both the ionic bond and the covalent bond. Within macromolecules such as proteins and nucleic acids, it can exist between two parts of the same molecule, and figures as an important constraint on such molecules' overall shape.

As the name "hydrogen bond" implies, one part of the bond involves a hydrogen atom. The hydrogen must be attached to a strongly electronegative heteroatom, such as oxygen, nitrogen or fluorine, which is called the hydrogen-bond donor. This electronegative element attracts the electron cloud from around the hydrogen nucleus and, by decentralizing the cloud, leaves the atom with a positive partial charge. Because of the small size of hydrogen relative to other atoms and molecules, the resulting charge, though only partial, nevertheless represents a large charge density. A hydrogen bond results when this strong positive charge density attracts a lone pair of electrons on another heteroatom, which becomes the hydrogen-bond acceptor.

The hydrogen bond is not like a simple attraction between point charges, however. It possesses some degree of orientational preference, and can be shown to have some of the characteristics of a covalent bond. This covalency tends to be more extreme when acceptors bind hydrogens from more electronegative donors.

Strong covalency in a hydrogen bond raises the questions: "To which molecule or atom does the hydrogen nucleus belong?" and "Which should be labelled 'donor' and which 'acceptor'? According to chemical convention, the donor generally is that atom to which, on separation of donor and acceptor, the retention of the hydrogen nucleus (or proton) would cause no increase in the atom's positive charge. The acceptor meanwhile is the atom or molecule that would become more positive by retaining the positively charged proton. Liquids that display hydrogen bonding are called associated liquids.

Hydrogen bonds can vary in strength from very weak (1-2 kJ mol⁻¹) to so strong (40 kJ mol⁻¹) so as to be indistinguishable from a covalent bond, as in the ion HF₂⁻. The length of hydrogen bonds depends on bond strength, temperature and pressure. The typical length of a hydrogen bond in water is 1.97 Å.

Hydrogen bond in water

The most ubiquitous, and perhaps simplest, example of a hydrogen bond is found between water molecules. In a discrete water molecule, water has two hydrogen atoms and one oxygen atom. Two molecules of water can form a hydrogen bond between them. The oxygen of one water molecule has two lone pairs of electrons, each of which can form a hydrogen bond with hydrogens on two other water molecules. This can repeat so that every water molecule is H-bonded with four other molecules (two through its two lone pairs, and two through its two hydrogen atoms.)

H-O-H^...O-H₂

Liquid water's high boiling point is due to the high number of hydrogen bonds each molecule can have relative to its low molecular mass. Water is unique because its oxygen atom has two lone pairs and two hydrogen atoms, meaning that the total number of bonds of a water molecule is four. (For example, hydrogen bromide- which has two lone pairs on the Br atom but only one H atom - can have a total of only two bonds.)

H-Br^...H-Br^...H-Br

In ice, the crystalline lattice is dominated by a regular array of hydrogen bonds which space the water molecules farther apart than they are in liquid water. This accounts for water's decrease in density upon freezing. In other words, the presence of hydrogen bonds enables ice to float, because this spacing causes ice to be less dense than liquid water.

Were the bond strengths more equivalent, one might instead find the atoms of two interacting water molecules partitioned into two polyatomic ions of opposite charge, specifically hydroxide and hydronium.(Hydronium ions are also known as 'hydroxonium' ions).

H-O⁻ H₃O⁺

Indeed, in pure water under conditions of standard temperature and pressure, this latter formulation is applicable only rarely; on average about one in every 5.5 * 10⁸ molecules gives up a proton to another water molecule, in accordance with the value of the dissociation constant for water under such conditions.

Hydrogen bond in proteins

Hydrogen bonding also plays an important role in determining the three-dimensional structures adopted by proteins and nucleic acids. In these macromolecules, bonding between parts of the same macromolecule cause it to fold into a specific shape, which helps determine the molecule's physiological or biochemical role. The double helical structure of DNA, for example, is due largely to hydrogen bonding between the base pairs, which link one complementary strand to the other and enable replication.

In proteins, hydrogen bonds form between the backbone oxygens and amide hydrogens. When the spacing of the amino acid residues participating in a hydrogen bond occurs regularly between positions i and i + 4, an alpha helix is formed. When the spacing is less, between positions i and i + 3, then a 3₁₀ helix is formed. When two strands are joined by hydrogen bonds involving alternating residues on each participating strand, a beta sheet is formed. (See also protein folding).

Dihydrogen bond

In a very recent development, hydrogen bonds have been noted between two hydrogen atoms having opposite polarity. An example occurs in the molecule H₃NBH₃ where the hydrogen atoms on nitrogen have a partial positive charge and the hydrogen atoms on boron have a partial negative charge. The resulting BH^...HN attractions cause the molecule to be a solid material rather than a gas as is the case in the closely related substance, H₃CCH₃. Because two hydrogen atoms are involved, this is termed a dihydrogen bond. See references below (Crabtree, et al.).

Symmetric hydrogen bond

Symmetric hydrogen bonds have been observed recently spectroscopically in formic acid at high pressure (>GPa). Each hydrogen atom forms a partial covalent bond with two atoms rather than one. Symmetric hydrogen bonds have been postulated in ice at high pressure (ice-X). See references below (Goncharov, et al.)

Advanced theory of the hydrogen bond

The hydrogen bond remains a fairly mysterious object in the theoretical study of quantum chemistry and physics. Most generally, the hydrogen bond can be viewed as a metric dependent electrostatic scalar field between two or more intermolecular bonds. This is slightly different than the intramolecular bound states of, for example, covalent or ionic bonds; however, hydrogen bonding is generally still a bound state phenomenon, since the interaction energy has a net negative sum. The question of the relationship between the covalent bond and the hydrogen bond remains largely unsettled, though the initial theory proposed by Linus Pauling suggests that the hydrogen bond has a partial covalent nature. While a lot of experimental data has been recovered for hydrogen bonds in water, for example, that provide good resolution on the scale of intermolecular distances and molecular thermodynamics, the kinetic and dynamical properties of the hydrogen bond in dynamical systems remains largely mysterious.

References

ISBN links support NWE through referral fees

• A New Intermolecular Interaction: Unconventional Hydrogen Bonds with Element-Hydride Bonds as Proton Acceptor Robert H. Crabtree, Per E. M. Siegbahn, Odile Eisenstein, Arnold L. Rheingold, and Thomas F. Koetzle Acc. Chem. Res. 1996, 29(7), 348 - 354.

• Polymerization of Formic Acid under High Pressure Alexander F. Goncharov, M. Riad Manaa, Joseph M. Zaug, Richard H. Gee, Laurence E. Fried, and Wren B. Montgomery Phys. Rev. Lett. 2005, 94, 065505.

ca:Pont d'hidrogen de:Wasserstoffbrückenbindung es:Enlace de hidrógeno fr:Liaison hydrogène it:Legame idrogeno he:קשרי מימן nl:Waterstofbrug ja:水素結合 nn:Hydrogenbinding pl:Wiązanie wodorowe ru:Водородная связь sv:Vätebindning

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The electrons in the MO of a bond are 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.

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. However, more complicated compounds such as metal complexes cannot be described by valence bond theory, and we need quantum chemistry (based on quantum mechanics) to help us understand these molecules.

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 us conceptualize the MO'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 bonds 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 MOs' structures and energies based on the AOs of the atoms they came from. Unlike pure ionic bonds, covalent bonds may have directed anisotropic properties.

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”.

These chemical bonds are intramolecular forces that keep atoms held together in molecules. There are also intermolecular forces that cause molecules to be attracted or repulsed by each other. These forces include ionic interactions, hydrogen bonds, dipole-dipole interactions, and induced dipole interactions.

Linus Pauling's book The Nature of the Chemical Bond is perhaps the most influential book on chemistry ever published.

References

• W. Locke (1997). Introduction to Molecular Orbital Theory. Retrieved May 18, 2005.
• Carl R. Nave (2005). HyperPhysics. Retrieved May 18, 2005.

External links

• Covalent Bonds and Molecular Structure

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