Metallic bond is a type of chemical bond that arises from electrostatic attraction between the conduction of electrons (in the form of a delocalized electron electron cloud) and a positively charged metal ion. This can be described as the division of free electrons between the lattice of positively charged ions (cations). Metal bond accounts for many of the physical properties of metals, such as strength, ductility, thermal and electrical resistivity and conductivity, opacity, and luster.
Metal bonding is not the only kind of chemical bond that metals can show, even as pure substances. For example, the gallium element consists of pairs that are covalently bonded in liquid and solid form - these pairs form a crystal lattice with metal bonds between the two. Another example of the metal-metal covalent bond is the flammable ion ( Hg 2
2 ).
Video Metallic bonding
History
When chemistry develops into science, it becomes clear that the metal forms most of the periodic table of elements and great progress is made in the description of salts that can be formed in reaction with acids. With the advent of electrochemistry, it became clear that the metal generally enters into the solution as positively charged ions and the oxidation reaction of the metal becomes well understood in the electrochemical series. An image emerges from the metal as a positive ion united by a sea of ​​negative electrons.
With the advent of quantum mechanics, this image is given a more formal interpretation in the form of a free electron model and its further expansion, an almost free electron model. In both these models, the electrons are seen as a gas moving through a solid lattice with energy that is essentially isotropic in that it depends on the magnitude of the square, not the direction of the momentum vector k . In a three-dimensional space, the set of dots of the highest level filled (the Fermi surface) should be a sphere. In almost free corrections of the model, the box-like Brillouin zone is added to k-space by the periodic potential experienced by the lattice (ionic), thus breaking down the isotropy slightly.
The emergence of X-ray diffraction and thermal analysis makes it possible to study the structure of crystalline solids, including metals and alloys, and the construction of phase diagrams becomes accessible. Despite all these advances, the nature of intermetallic compounds and alloys remains largely a mystery and their studies are often empirical. Chemists are generally away from everything that does not seem to follow Dalton's law in double proportion and the matter is considered a domain of different science, metallurgy.
The almost-free electron model is enthusiastically picked up by several researchers in this field, notably Hume-Rothery, in an attempt to explain why certain intermetallic alloys with certain compositions will be formed and others not. Initially his business was quite successful. The idea is to add electrons to inflate the balloon Fermi ball inside the Brillouin-box series and determine when a particular box will be full. It does predict the alloy composition observed quite a lot. Unfortunately, as soon as the cyclotron resonance becomes available and the shape of the balloon can be determined, it is found that the assumption that the balloon is round can not stand at all, except perhaps in the case of cesium. This reduces many conclusions for an example of how a model can sometimes provide a whole set of correct predictions, but it is still wrong.
The free electron catastrophe shows researchers that the model assuming that ions are in a sea of ​​free electrons requires modification, and a number of models of quantum mechanics such as the calculation of a band structure based on a molecular orbital or density functional theory developed. In this model, one either departs from a neutral atomic atomic orbital that shares their electrons or (in the case of density functional theory) departs from the total electron density. The picture of free electrons, however, remains dominant in education.
The electronic ribbon structure model became the main focus not only for metal studies but more than that for semiconductor studies. Along with the electronic state, the vibration status is also shown to form the band. Rudolf Peierls points out that, in the case of a one-dimensional array of metallic atoms, say hydrogen, instability must emerge which would lead to the breaking of such chains into individual molecules. This sparked an interest in the general question: When is the collective metal bond stable and when does the more localized bond form replace its place? Much research goes into the study of grouping of metal atoms.
As strong as the concept of band structure is evident in the description of metal bonding, it does have its drawbacks. It still approaches one electron to many multi-body problems. In other words, the energy states of each electron are described as if all other electrons are just forming a homogeneous background. Researchers like Mott and Hubbard recognize that this might be suitable for the delocalized electron s- and p, but for d-electrons, and even more for f-electrons, the interaction with electrons (and atomic displacements) in the local environment can be stronger than delocalization that leads to big bands. Thus, the transition of unpaired electrons that are localized to electrons grouped in metal bonds becomes more easily understood.
Maps Metallic bonding
The properties of the metal bond
The combination of two phenomena produces a metal bond: the delocalization of electrons and the availability of a much larger amount of energy states delocalized than delocalized electrons. The latter can be called electron deficiency.
In 2D
Graphene is an example of a two-dimensional metal bond. The metal bond is similar to the aromatic bond in benzene, naphthalene, anthracene, ovalena, and so on.
In 3D
Metal aromaticity in metal groups is another example of delocalisation, this time often in three-dimensional entities. The metal takes the principle of delocalisation to the extreme and it can be said that the crystals of the metal represent a molecule in which all conductive electrons are delocalized in all three dimensions. This means that in a person's metal is generally unable to distinguish the molecule, so the metal bond is neither intra-nor intermolecular. 'Nonmolecular' would probably be a better term. Metallic bonds are mostly non-polar, because even in alloys there is little difference between the electronegativity of atoms participating in bonding interactions (and, in pure elemental metals, none at all). Thus, metal bonding is a highly delocalized form of covalent bonding communal. In a sense, metal bonding is not a 'new' type of bond at all, therefore, and it describes bonds only as present in the pieces of the condensed matter, whether crystal, liquid, or even glass solid. Metallic vapors with contrast are often atoms (Hg) or sometimes contain molecules like Na 2 that are united by more conventional covalent bonds. This is why it is not right to talk about a 'metal bond'.
The most prominent delocalizations for s - and p -electrons. For cesium it is so strong that the electrons are almost free of cesium atoms to form a gas that is limited only by the metal surface. For cesium, therefore, the image of the Cs ion held together by the negatively charged electron gas is not very inaccurate. For other elements less free electrons, because they still experience the potential of metal atoms, sometimes quite strong. They require the maintenance of more complex quantum mechanics (for example, tight binding) in which atoms are viewed as neutral, such as carbon atoms in benzene. For d - and especially f -electrons, delocalisation is not strong at all and this explains why these electrons can continue to behave as unpaired electrons that retain their spin, adding interesting magnetic properties for this metal.
Disadvantages and mobility of electrons
Metal atoms contain several electrons in their valence shells relative to their period or energy level. They are elements that lack electrons and communal divisions do not change that. There are still more available energy states than the shared electrons. Therefore, both requirements for conductivity are met: strong desocalisation and partially charged energy bands. Such electrons can easily change from one energy state to a slightly different one. Thus, not only do they become delocalized, forming a sea of ​​electrons that penetrate the lattice, but they can also migrate through the lattice when an external electric field is imposed, leading to electrical conductivity. Without the terrain, there are electrons that move equally in all directions. Under the field, some will slightly adjust their circumstances, adopting different wave vectors. As a result, there will be more movement in one direction than the other and a net current will be generated.
The freedom of conduction electrons to migrate also gives the metal atoms, or layers, the capacity to slide past each other. Locally, bonds can be easily damaged and replaced with new ones after deformation. This process does not affect the bonds of the communal metal very much. This gives rise to a typical phenomenon characteristic of elasticity and ductility of metals. This is especially true for pure elements. In the presence of soluble discharges, defects in the lattice that serve as splitting points can be blocked and the material becomes harder. Gold, for example, is very soft in its pure form (24-karat), which is why an 18-carat or lower alloy is preferred in jewelry.
Metal is also a good conductor of heat, but electron conduction only contributes in part to this phenomenon. The collective vibration (ie, delocalized) atoms known as phonons that travel through solid as waves, contribute strongly.
However, the latter also applies to substances such as diamonds. It does heat well enough but no electricity. The latter is not a consequence of the fact that delocalisation does not exist in diamonds, but only carbon that lacks electrons. Electron deficiency is an important point in distinguishing metal from more conventional covalent bonds. Thus, we must alter the expression given above to be: The metal bond is a highly delocalized communal form of the electron-deficient covalent bond .
Metallic radius
The metallic radius is defined as one-half the distance between two adjacent metal ions in the metallic lattice. The radius depends on the nature of the atom and its environment - in particular, the coordination number (CN), which in turn depends on the temperature and pressure applied.
When comparing periodic trends in atomic sizes, it is often desirable to apply so-called Goldschmidt correction, which converts radii to atomic values ​​will have if they are 12-coordinated. Since the metal radius is always largest for the highest coordination number, the correction for the less dense coordination involves multiplying by x, where 0 & lt; x & lt; 1. Specifically, for CN = 4, x = 0.88; for CN = 6, x = 0.96, and for CN = 8, x = 0.97. This correction is named after Victor Goldschmidt who obtained the numerical values ​​quoted above.
The fingers follow a general periodic trend: they decrease throughout the period due to an increase in effective nuclear charge, which is not matched by an increase in the number of valence electrons. The radius also rises down the group due to the increase of the major quantum numbers. Between row 3 and 4, lanthanide contractions are observed - there is a slight increase in radius down the group due to the poorly protected f orbitals.
Strength ties
The atoms in the metal have strong tensile strength between them. A lot of energy is needed to overcome it. Therefore, the metal often has a high boiling point, with tungsten (5828 K) becoming very high. The exceptional exception are the elements of the zinc group: Zn, Cd, and Hg. Their electron configuration ends at... ns 2 and this appears to resemble a noble gas configuration such as helium that gets more and more when it drops off in the periodic table because the energy distance to the empty np orbitals becomes larger. Therefore, the metal is relatively volatile, and is avoided in ultra-high vacuum systems.
Otherwise, the metal bond can be very strong, even in molten metal, like Gallium. Although the gallium will melt from the heat of one's hands just above room temperature, its boiling point is not far from copper. Thus, gallium gallium is a very non-flammable liquid thanks to a strong metal bond.
The strong bond of the metal in liquid form shows that the metal bonding energy is not a strong function of the metal bonding direction; the lack of attachment of these bonds is a direct consequence of the delocalization of electrons, and is best understood in contrast to the directional bond of the covalent bond. The energy of the metal bond is thus largely a function of the number of electrons that surround the metal atom, as exemplified by the embedded atom model. This usually produces metals assuming relatively simple and closed crystal structures, such as FCC, BCC, and HCP.
With a high enough cooling rate and an appropriate alloying composition, metal bonds can occur even in glasses with amorphous structures.
Many biochemicals are mediated by weak interactions of metal ions and biomolecules. These interactions and associated conformational changes have been measured using multiple polarization interferometry.
The solubility and compound formation
Metals are not soluble in water or organic solvents unless they experience reactions with them. Usually this is an oxidation reaction that robs the metal atoms of their peripheral electrons, destroying metal bonds. But metals often readily dissolve each other while retaining their metal bonding characters. Gold, for example, is easily soluble in mercury, even at room temperature. Even in solid metal, its solubility can be very wide. If the structure of the two metals is the same, there is even complete solid solubility, as in the case of electrons, silver and gold alloys. Sometimes, however, two metals will form alloys with different structures from one of two parents. One could call these materials as metal compounds, but, since materials with metal bonds are not usually molecular, Dalton's law of integral proportions is invalid and often various stoichiometric ratios can be achieved. It is better to abandon concepts such as 'pure substances' or 'solutes' are such cases and talk about phases instead. The study of these phases has traditionally been more a metallurgical domain than a chemical, although the two fields overlap.
Localization and grouping: from bond to
The bonding of metals in complex compounds does not necessarily involve all the constituent elements evenly. It is possible to have elements or more that do not take part at all. One can imagine conducting electrons flowing around them like a river around an island or a large rock. It is possible to observe which element takes part, for example, by looking at the core level in the X-ray photoelectron spectroscopy (XPS) spectrum. If an element takes part, its peak tends to tilt.
Some intermetallic materials for example do exhibit metal clusters, reminiscent of molecules and these compounds are more of a chemical topic than metallurgy. The formation of clusters can be seen as a way to 'condense' (localize) the bond of electron deficiency into a more localized bond. Hydrogen is an extreme example of this form of condensation. At high pressure it is metal. The nucleus of the planet Jupiter can be said to be held together by a combination of metal bonds and high pressure caused by gravity. At lower pressures the bonds become fully localized into ordinary covalent bonds. This localization is so complete that the more well known (more than> 2 ) gas H results. A similar argument applies to elements such as boron. Although it lacks electrons compared to carbon, it does not form metal. Instead, it has a number of complex structures in which the icosahedral clause B 12 dominates. Filling the density wave is a related phenomenon.
When these phenomena involve the movement of atoms toward or away from each other, they can be interpreted as merging between electronic and vibrational status (ie phonons) of material. Different electron-phonon interactions are thought to cause very different results at low temperatures, ie, superconductivity. Instead of blocking the mobility of charge carriers by forming electron pairs in local bonds, Cooper pairs are formed that no longer have resistance to their mobility.
Optical properties
The presence of oceans of cellular carriers has a profound effect on the optical properties of metals. They can only be understood by considering electrons as collective rather than considering the state of individual electrons involved in more conventional covalent bonds.
Light consists of a combination of electric and magnetic fields. The electric field is usually capable of generating the elastic response of the electrons involved in metal bonding. The result is that photons can not penetrate very far into the metal and are usually reflected. They bounce off, though some can also be absorbed. This applies equally to all photons of the visible spectrum, which is why metals are often white or grayish with speculative reflections typical of metallic luster. The balance between reflection and absorption determines how white or how gray they are, even if the surface stains can obscure the observations. Silver, a very good metal with high conductivity is one of the most white.
The main exceptions are reddish copper and yellowish gold. The reason for their color is that there is an upper limit of the frequency of light that the metal electrons can easily respond to, the plasmon frequency. At the plasmon frequency, the frequency dependent dielectric function of the free electron electron changes from negative (reflects) to positive (transmission); Higher frequency photons are not reflected on the surface, and do not contribute to metal colors. There are some ingredients like indium tin oxide (ITO) which are metal conductors (actually semiconductor degeneration) where these thresholds are in infrared, that is why they are transparent in visible mirrors, but are good in IR.
For silver, the limiting frequency is in the distant UV, but for copper and gold closer to the visible. This explains the colors of these two metals. On the surface of the metal resonance effect known as plasmons surface may occur. They are collective oscillations of conduction electrons like ripples in the ocean of electronics. However, even if the photons have enough energy they usually do not have enough momentum to set the ripple to move. Therefore, plasmons are difficult to exclude on large metals. This is why gold and copper still look like sparkling metal though with a bit of color. However, in colloidal gold the metal bond is confined to small metal particles, preventing the plasmon oscillation wave from 'escape'. Hence the momentum selection rule is broken, and plasmon resonance causes a very intense absorption in green with a beautiful purple-red color produced. Such colors are a stronger order of magnitude than normal perceptions seen in dyes and the like that involve individual electrons and their energy status.
See also
- Bond in solid form
- Metal aromaticity
References
Source of the article : Wikipedia
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