examples of metallic bonds
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Metallic bonding is the electromagnetic interaction between delocalized electrons, called conduction electrons, 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.
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 (Hg|2|2+), which forms covalent metal-metal bonds.
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
An ionic bond is a type of chemical bond that involves a metal and a nonmetalion (or polyatomic ions such as ammonium) through electrostatic attraction. In short, it is a bond formed by the attraction between two oppositely charged ions.
The metal donates one or more electrons, forming a positively charged ion or cation with a stable electron configuration. These electrons then enter the non metal, causing it to form a negatively charged ion or anion which also has a stable electron configuration. The electrostatic attraction between the oppositely charged ions causes them to come together and form a bond.
For example, common table salt is sodium chloride. When sodium (Na) and chlorine (Cl) are combined, the sodium atoms each lose an electron, forming cations (Na+), and the chlorine atoms each gain an electron to form anions (Clâˆ’). These ions are then attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).
- Na + Cl â†’ Na+ + Clâˆ’â†’ NaCl
The removal of electrons from the atoms is endothermic and causes the ions to have a higher energy. There may also be energy changes associated with breaking of existing bonds or the addition of more than one electron to form anions. However, the attraction of the ions to each other lowers their energy. Ionic bonding will occur only if the overall energy change for the reaction is favourable â€“ when the bonded atoms have a lower energy than the free ones. The larger the resulting energy change the stronger the bond. The low electronegativity of metals and high electronegativity of non-metals means that the energy change of the reaction is most favorable when metals lose electrons and non-metals gain electrons.
Pure ionic bonding is not known to exist. All ionic compounds have a degree of covalent bonding. The larger the difference in electronegativity between two atoms, the more ionic the bond. Ionic compounds conduct electricity when molten or in solution. They generally have a high melting point and tend to be soluble in water.
Ionic compounds in the solid state form lattice structures. The two principal factors in determining the form of the lattice are the relative charges of the ions and their relative sizes. Some structures are adopted by a number of compounds; for example, the structure of the rock salt sodium chloride is also adopted by many alkali halides, and binary oxides such as MgO.
Strength of an ionic bond
For a solid crystalline ionic compound the enthalpy change in forming the solid from gaseous ions is termed the lattice energy. The experimental value for the lattice energy can be determined using the Born-Haber cycle. It can also be calculated using the Born-LandÃ© equation as the sum of the electrostatic potential energy, calculated by summing interactions between cations and anions, and a short range repulsive potential energy term. The electrostatic potential can be expressed in terms of the inter-ionic separation and a constant (Madelung constant) that takes account of the geometry of the crystal. The Born-LandÃ© equation gives a reasonable fit to the lattice energy of e.g. sodium chloride where the calculated value is âˆ’756 kJ/mol which compares to âˆ’787 kJ/mol using the Born-Haber cycle.
Ions in crystal lattices of purely ionic compounds are spherical; however, if the positive ion is small and/or highly charged, it will distort the electron cloud of the negative ion, an effect summarised in Fajans' rules. This polarization of the negative ion leads to a build-up of extra charge density between the two nuclei, i.e., to partial covalency. Larger negative ions are more easily polarized, but the effect is usually only important when positive ions with charges of 3+ (e.g., Al3+) are involved. However, 2+ ions (Be2+) or even 1+ (Li+) show some polarizing power because their sizes are so small (e.g., LiI is ionic but has some covalent bonding present). Note that this is not the ionic polarization effect which refers to displacement of ions in the lattice due to the application of an electric field.
Ionic versus covalent bonds
In an ionic bond, the atoms are bound by attraction of opposite ions, whereas, in a covalent bond, atoms are bound by sharing electrons. In covalent bonding, the molecular geometry around each atom is determined by VSEPR rules, whereas, in ionic materials, the geometry follows maximum packing rules.
In reality, purely ionic bonds do no
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Answers:draw circles packed together without the bonds showing just sticked together off something like this http://farm1.static.flickr.com/19/161217646_dcd577049d.jpg?v=0 http://images.google.com/imgres?imgurl=http://media.tiscali.co.uk/images/feeds/hutchinson/ency/0013n055.jpg&imgrefurl=http://www.talktalk.co.uk/reference/encyclopaedia/hutchinson/m0030538.html&usg=__yGSgiIZFdQxm66jW7fcYYNM-yrc=&h=186&w=382&sz=13&hl=en&start=87&um=1&itbs=1&tbnid=-cbwNXuq9JeHzM:&tbnh=60&tbnw=123&prev=/images%3Fq%3Dmetallic%2Bbonding%26start%3D72%26um%3D1%26hl%3Den%26client%3Dfirefox-a%26sa%3DN%26rls%3Dorg.mozilla:en-US:official%26ndsp%3D18%26tbs%3Disch:1 http://media.tiscali.co.uk/images/feeds/hutchinson/ency/0013n055.jpg
Answers:Alright: 1) Metallic bonds are weak bonds that metallic atoms form to temporarily share electrons. Zn2 is a good example. Zn atoms will bond together for short periods of time, until a highly electronegative atom comes to take the electrons each atom has. 2) Covalent bonds are bonds that involve atoms sharing electrons equally (non-polar) and unequally (polar). Such examples would be hypochlorite, ClO, and water, H20. The elctronegativity difference in ClO is just about 0, and in water it is about 1.3. These indicate the polarity of the molecule. 3) Ionic bonds are bonds that involve one weak electronegative atom and one strong electronegative atom. Such examples would be NaCl, NaF, KCl, KF, CaCl2, etc. The electronegativity difference is large (usually >2.0). Hope this helps. If you need more examples, see if Wikipedia can help.
Answers:From Wikipedia: Although the term metallic bond is often used in contrast to the term covalent bond, it is more preferable to use the term metallic bonding, because this type of bonding is collective in nature and a single "metallic bond" does not exist. Metallic bonding is the electromagnetic interaction between delocalized electrons, called conduction electrons, and the metallic nuclei within metals. I don't understand why you call this metallic bonding equation. I assume that it is an equation for the reactivity of metals. Metal Reactivity: http://en.wikipedia.org/wiki/Activity_series_of_metals You realise that Potassium is more reactive than copper so copper will not react with potassium nitrate as it sort of binds more to the nitrate anion However, iron is more reactive than copper, so it will react with copper(II)nitrate and gives you -> Iron(II) nitrate (aq) + Copper (s) 2 Li + 2 H2O 2 LiOH + H2 Lithium is more reactive than Hydrogen so it will react with H20. It is an alkaline metal (Grp I in periodic table) Hence it will form an alkaline --> Hydroxide. And you have learnt that more reactive metals liberates Hydrogen gas with water.
Answers:Wow, these people are giving the worst answers. So I don't fully understand how metallic bonds work, but they're similar in that the bonded atoms don't share any electrons. In an ionic bond, some of the electrons from one atom movie into the outermost shell of the other atom, so the charges become opposite but equal. For example in NaCl, an electron from the sodium goes moves into the chlorine, making the Na+ and Cl-. Because the charges are opposite, the atoms are attracted to each other. And for the covalent bond question, the more bonds there are, the stronger it is.