Metallic bond

Metallic bonding is the electromagnetic interaction between delocalized electrons, called conduction electrons and gathered in an "electron sea", and the metallic nuclei within metals. Understood as the sharing of "free" electrons among a lattice of positively-charged ions (cations), metallic bonding is sometimes compared with that of molten salts; however, this simplistic view holds true for very few metals. In a more quantum-mechanical view, the conduction electrons divide their density equally over all atoms that function as neutral (non-charged) entities. Metallic bonding accounts for many physical properties of metals, such as strength, malleability, ductility, thermal and electrical conductivity, opacity, and luster.[1][2][3][4]

Although the term "metallic bond" is often used in contrast to the term "covalent bond", it is preferable to use the term metallic bonding, because this type of bonding is collective in nature and a single "metallic bond" does not exist. Not all metals exhibit metallic bonding: one such example is the mercurous ion (Hg2+2), which forms covalent metal-metal bonds.

History

The nature of metals has fascinated humankind for many centuries, because these materials provided people with tools of unsurpassed properties both in war and in peace. The reason for their properties and the nature of the bonding that keeps them together remained a mystery for centuries, even though great progress was made in their preparation and processing.

As chemistry developed into a science it became clear that metals formed the large majority of the periodic table of the elements and great progress was made in the description of the salts that can be formed in reactions with acids. With the advent of electrochemistry it became clear that metals generally go into solution as positively charged ions and the oxidation reactions of the metals became well understood in the electrochemical series. A picture emerged of metals as positive ions held together by an ocean of negative electrons.

With the advent of quantum mechanics this picture was given more formal interpretation in the form of the free electron model and its further extension, the nearly-free electron model. In both of these models the electrons are seen as a gas traveling through the lattice of the solid with an energy that is essentially isotropic in that it depends on the square of the magnitude, not the direction of the momentum vector k. In three-dimensional k-space, the set of points of the highest filled levels (the Fermi surface) should therefore be a sphere. In the nearly-free correction of the model, box-like Brillouin zones are added to k-space by the periodic potential experienced from the (ionic) lattice.

The advent of X-ray diffraction and thermal analysis (initially DTA) made it possible to study the structure of crystalline solids, including metals and their alloys, and the construction of phase diagrams became accessible. Despite all this progress the nature of intermetallic compounds and alloys largely remained a mystery and their study was often empirical. Chemists generally steered away from anything that did not seem to follow Dalton's laws of multiple proportions and the problem was considered the domain of a different science, metallurgy.

The almost-free electron model was eagerly taken up by some researchers in this field, notably Hume-Rothery in an attempt to explain why certain intermetallic alloys with certain compositions would form and others would not. Initially his attempts were quite successful. Basically his idea was to add electrons to inflate the spherical Fermi-balloon inside the series of Brillouin-boxes and determine when a certain box would be full. This indeed predicted a fairly large number of observed alloy compositions. Unfortunately, as soon as cyclotron resonance became available and the shape of the balloon could be determined, it was found that the assumption that the balloon was spherical did not hold at all, except perhaps in the case of caesium. This reduced many of the conclusions to examples of how a wrong model can sometimes give a whole series of correct predictions.

The free-electron debacle showed researchers that the model which assumed the ions were in a sea of free electrons needed modification, and a number of quantum mechanical models such as band structure calculations based on molecular orbitals or the density functional theory were developed. In these models one departs either from the atomic orbitals of neutral atoms that share their electrons or (in the case of DFT) departs from the total electron density. The free-electron picture has nevertheless remained a dominant one in education.

The electronic band structure model became a major focus not only for the study of metals, but even more so for the study of semiconductors. Together with the electronic states, the vibrational states were also shown to form bands. Peierls showed that in the case of a one-dimensional row of metallic atoms, say hydrogen, an instability had to arise that would lead to the break up of such a chain into individual molecules. This sparked an interest in the general question: when is collective metallic bonding stable and when will a more localized form of bonding take its place? Much research went into the study of clustering of metal atoms.

As powerful as the concept of the band structure proved to be in the description of metallic bonding it does have a drawback. It remains a one-electron approximation to a multitudinous many-body problem. In other words, the energy states of each electron are described as if all the other electrons simply form a homogeneous background. Researchers like Mott and Hubbard realized that this was perhaps appropriate for strongly delocalized s- and p-electrons but for d-electrons, and even more for f-electrons the interaction with electrons (and atomic displacements) in the local environment may become stronger than the delocalization that leads to broad bands. Thus the transition from localized unpaired electrons to itinerant ones partaking in metallic bonding became more comprehensible.
The nature of metallic bonding

The combination of two phenomena gives rise to metallic bonding: delocalization of electrons and the availability of a far larger number of delocalized energy states than of delocalized electrons. The latter could be called electron deficiency.

Delocalization
In 2D

Delocalization—bonding that involves more than one pair of atoms held together by one pair of electrons—is most familiar from the example of benzene C6H6, where six electrons from six carbon atoms are engaged in joint aromatic bonding. However, there are other examples like the three-center two-electron bond, prevalent in boron chemistry. The principle can easily be extended over larger aromatic molecules like naphthalene, anthracene and if the process is taken to its extreme: graphite. The latter is an example of a system delocalized in two dimensions. Interestingly, there is an isoelectronic analog of benzene, B3N3H6 (borazine) for which the same argument holds. It has very similar properties to benzene.[5] When extended indefinitely, hexagonal boron nitride, BN, is obtained, with a structure identical to that of graphite apart from the alternation between boron and nitrogen in each ring. This material is a semiconductor, exemplifying that delocalization is a necessary but not sufficient requirement for conductivity. Electrical conductivity does occur in graphite, because the π and π*-like bands overlap, making it a semimetal, with partly filled bands, fulfilling the other requirement for conductivity.

In 3D

Metal aromaticity in metal clusters is another example of delocalization, this time often in three-dimensional entities. Metals take the delocalization principle to its extreme and one could say that a crystal of a metal represents a single molecule over which all conduction electrons are delocalized in all three dimensions. This means that inside the metal one can generally not distinguish molecules so that the metallic bonding is neither intra- nor intermolecular. 'Nonmolecular' would perhaps be a better term. Metallic bonding is mostly non-polar, because even in alloys there is little difference among the electronegativities of the atoms participating in the bonding interaction (and in pure elemental metals, none at all). Thus metallic bonding is an extremely delocalized communal form of covalent bonding. In a sense metallic bonding is not a 'new' type of bonding at all therefore and it only describes the bonding as present in a chunk of condensed matter, be it crystalline solid, liquid or even glass. Metallic vapors by contrast are often atomic (Hg) or at times contain molecules like Na2 held together by a more conventional covalent bond. This is why it is not correct to speak of a single 'metallic bond'.

The delocalization is most pronounced for s- and p-electrons. For caesium it is so strong that the electrons are virtually free from the caesium atoms to form a gas only constrained by the surface of the metal. For caesium therefore the picture of Cs+-ions held together by a negatively charged electron gas is not too inaccurate [6]. For other elements the electrons are less free, in that they still experience the potential of the metal atoms, sometimes quite strongly. They require a more intricate quantum mechanical treatment (e.g. tight binding) in which the atoms are viewed as neutral much like the carbon atoms in benzene. For d- and especially f-electrons the delocalization is not strong at all and this explains why these electrons are able to continue behaving as unpaired electrons that retain their spin, adding interesting magnetic properties to these metals.

Electron deficiency and mobility

Metal atoms contain few electrons in their valence shells relative to their periods or energy levels. They are electron deficient elements and the communal sharing does not change that. There remain far more available energy states than there are shared electrons. Both requirements for conductivity are therefore fulfilled: strong delocalization and partly filled energy bands. Such electrons can therefore easily change from one energy state into a slightly different one. Thus, not only do they become delocalized, forming a sea of electrons permeating the lattice, but they are also able to migrate through the lattice when an external electrical field is imposed, leading to electrical conductivity. Without the field there are electrons moving equally in all directions. Under the field some will adjust their state slightly, adopting a different wave vector. Consequently, there will be more moving one way than the other and a net current will result.

The freedom of conduction electrons to migrate also gives metal atoms, or layers of them, the capacity to slide past each other. Locally bonds can easily be broken and replaced by new ones after the deformation. This process does not affect the communal metallic bonding very much. This gives rise to metals' typical characteristic phenomena of malleability and ductility. This is particularly true for pure elements. In the presence of dissolved impurities the defects in the lattice that function as cleavage points may get blocked and the material becomes harder. Gold for example is very soft in pure form (24 kt), which is why for jewellery alloys of 18 kt or lower are preferred.

Metals are typically also good conductors of heat, but the conduction electrons only contribute partly to this phenomenon. Collective (i.e. delocalized) vibrations of the atoms known as phonons that travel through the solid as a wave, contribute strongly.

However, the latter also holds for a substance like diamond. It conducts heat quite well but not electricity. The latter is not a consequence of the fact that delocalization is absent in diamond, but simply that carbon is not electron deficient . The position of carbon in the middle of its period in the Periodic Table makes that there are precisely enough electrons to fill the energy states. Under a field electrons are not able to adopt a different wave vector because there are no empty states to move into. This makes a current impossible in this wide band gap semiconductor. However, as soon as charge carriers are introduced by doping the crystal with a suitable impurity the resulting charge carriers are as mobile as in a metal, though far fewer in number. Even without doping the vibrational motions (the phonons) are delocalized over the crystal explaining the heat conduction. Still the bonding in diamond is better described as covalent than as metallic if only because there is a very strong directional preference for tetrahedral stacking, producing a structure that is extremely hard to deform and by no means close packed.

Clearly, the electron deficiency is an important point in distinguishing metallic from more conventional covalent bonding. Thus, we should amend the expression given above into:
Metallic bonding is an extremely delocalized communal form of electron deficient[7] covalent bonding.
Metallic radius

Metallic radius is defined as one half of the distance between neighbouring atoms in the metal solid. This radius depends on the nature of the atom as well as its environment, specifically on the coordination number which in turn depends on the temperature and applied pressure.

When comparing periodic trends in the size of atoms it is often desirable to apply so-called Goldschmidt correction which converts the radii to the values which the atoms would have if they were 12-coordinated. Since metallic radii are always biggest for the highest coordination number, correction for less dense coordinations involves dividing by x, where 0 < x < 1. Specifically, for cn. 4 x = 0.88, for cn. 6 x = 0.96, and for cn. 8 x = 0.97. The correction is named after Victor Goldschmidt who obtained the numerical values quoted above.[8]

The radii follow general periodic trends: they decrease across the period due to increase in the effective nuclear charge which is not offset by the increased number of valence electrons. The radii also decrease down the group due to incrase in principal quantum number. Between rows 3 and 4 the lanthanide contraction is observed - there is very little increase of the radius down the group due to the presence of poorly shielding f orbitals.
Strength of the bond

The atoms in metals have a strong attractive force between them. Much energy is required to overcome it. Therefore, metals often have high boiling points, with tungsten (5828 K) being extremely high. A remarkable exception are the elements of the zinc group: Zn, Cd,and Hg. Their electron configuration ends in ...ns2 and this comes to resemble a noble gas configuration like that of helium more and more when going down in the periodic table because the energy distance to the empty np orbitals becomes larger. These metals are therefore relatively volatile. Zinc for example is unsuitable for use in ultra high vacuum applications.

Otherwise, metallic bonding can be very strong, even in the melt. Gallium is a good example of that. Even though it melts by the heat of one's hand just above room temperature, its boiling point is not far from that of copper. Molten gallium is therefore a very nonvolatile liquid thanks to its strong metallic bonding.

The latter also exemplifies that metallic bonding due to its delocalization in all directions is often not very particular about the directionality of the bonding. There is typically a preference for close packing of the atoms, such as face or body centered cubic arrangements, but in the case of liquid gallium the stacking is not regular, at least not at long range and bond angles are easily changed.

Given high enough cooling rates and appropriate alloy composition metallic bonding can even occur in glasses with an amorphous structure.

Much biochemistry is mediated by the weak interaction of metal ions and biomolecules. Such interactions and their associated conformational change has been measured using dual polarisation interferometry.
Solubility and compound formation

Metals are insoluble in water or organic solvents unless they undergo a reaction with them. Typically this is an oxidation reaction that robs the metal atoms of their itinerant electrons, destroying the metallic bonding. However metals are often readily soluble in each other while retaining the metallic character of their bonding. Gold for example dissolves easily in mercury, even at room temperature. Even in solid metals the solubility can be extensive. If the structures of the two metals are the same there can even be complete solid solubility as in the case of electrum, the alloys of silver and gold. At times however two metals will form alloys with different structures than either of the two parents. One could call these materials metal compounds, but because materials with metallic bonding are typically not molecular, Dalton's law of integral proportions is not valid and often a range of stoichiometric ratios can be achieved. It is better to abandon such concepts as 'pure substance' or 'solute' is such cases and speak of phases instead. The study of such phases has traditionally been more the domain of metallurgy than of chemistry, although the two fields overlap considerably.
Localization and clustering: from bonding to bonds

The metallic bonding in complicated compounds does not necessarily involve all constituent elements equally. It is quite possible to have an element or more that do not partake at all. One could picture the conduction electrons flowing around them like a river around an island or a big rock. It is possible to observe which elements do partake, e.g. by looking at the core levels in an XPS spectrum. If an element partakes its peaks tend to be skewed.

Some intermetallic materials e.g. do exhibit metal clusters, reminiscent of molecules and these compounds are more a topic of chemistry than of metallurgy. The formation of the clusters could be seen as a way to 'condense out' (localize) the electron deficient bonding into bonds of a more localized nature. Hydrogen is an extreme example of this form of condensation. At high pressures it is a metal. The core of the planet Jupiter could be said to be held together by a combination of metallic bonding and high pressure induced by gravity. At lower pressures however the bonding becomes entirely localized into a regular covalent bond. The localization is so complete that the (more familiar) H2 gas results. A similar argument holds for an element like boron. Though it is electron deficient compared to carbon, it does not form a metal. Instead it has a number of complicated structures in which icosahedral B12 clusters dominate. Charge density waves are a related phenomenon.

As these phenomena involve the movement of the atoms towards or away from each other, they can be interpreted as the coupling between the electronic and the vibrational states (i.e. the phonons) of the material. A different such electron-phonon interaction is thought to cause a very different result at low temperatures, that of superconductivity. Rather than blocking the mobility of the charge carriers by forming electron pairs in localized bonds, Cooper-pairs are formed that no longer experience any resistance to their mobility.
Optical properties

The presence of a ocean of mobile charge carriers has profound effects on the optical properties of metals. They can only be understood by considering the electrons as a collective rather than considering the states of individual electrons involved in more conventional covalent bonds.

Light consists of a combination of an electrical and a magnetic field. The electrical field is usually able to excite an elastic response from the electrons involved in the metallic bonding. The result is that photons are not able to penetrate very far into the metal and are typically reflected. They bounce off, although some may also be absorbed. This holds equally for all photons of the visible spectrum, which is why metals are often silvery white or grayish with the characteristic specular reflection of metallic lustre. The balance between reflection and absorption determines how white or how gray they are, although surface tarnish can obscure such observations. Silver, a very good metal with high conductivity is one of the whitest.

Notable exceptions are reddish copper and yellowish gold. The reason for their color is that there is an upper limit to the frequency of the light that metallic electrons can readily respond to, it is known as the plasmon frequency. For light that oscillates much faster than this limit the material becomes transparent: the charge carriers are simply too sluggish to follow the rapidly oscillating photons and let them pass unhindered. There are some materials like indium tin oxide (ITO) that are metallic conductors (actually degenerate semiconductors) for which this threshold is in the infrared[9] which is why they are transparent in the visible, but good mirrors in the IR.

For silver the limiting frequency is in the far UV, but for copper and gold it is closer to the visible. This explains the colors of these two metals. At the surface of a metal resonance effects known as surface plasmons can result. They are collective oscillations of the conduction electrons like a ripple in the electronic ocean. However, even if photons have enough energy they usually do not have enough momentum to set the ripple in motion. Therefore, plasmons are hard to excite on a bulk metal. This is why gold and copper still look like lustrous metals albeit with a dash of color. However, in colloidal gold the metallic bonding is confined to a tiny metallic particle preventing the oscillation wave of the plasmon to 'run away'. The momentum selection rule is therefore broken and the plasmon resonance causes an extremely intense absorption in the green and a beautiful wine-red color. These colors are orders of magnitude more intense than ordinary absorptions seen in dyes and the like that involve individual electrons and their energy states.

Chemistry: What is a metal? (Metallic Bonds)