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scott_alexsk
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How are metallic bonds formed? Does it have anything to do with the merging of d orbitals to obtain a half full shell configuration or is it something else? Thank you for your time.
-Scott
-Scott
http://en.wikipedia.org/wiki/Transition_metals#Electronic_configurationTransition elements tend to have high tensile strength, density and melting and boiling points. As with many properties of transition metals, this is due to d orbital electrons' ability to delocalise within the metal lattice. In metallic substances, the more electrons shared between nuclei, the stronger the metal.
scott_alexsk said:Thank you for replying Aldriono and Astronuc. I really don't know much about metallic bonds except for that they conduct electricity. Do ionic bonds regularly conduct electricity?
-Scott
Gokul43201 said:One way to simplistically think about things is the molecular-orbital approach. When you overlap a pair of orbitals of equal energy you create two new energy levels (above and below the original energy) called bonding and antibonding orbitals.
First off, I said what followed was a simplistic approach to form an intuition for what's happening. That said, the above approach does accurately convey some of the key elements of the picture. While it's not a complete picture (for instance, so far, none of the features resulting from having a discrete translational symmetry are explained), it is certainly incorrect to say that this applies only to covalent bonds between a small number of atoms.photon79 said:But this is true only in the case of covalent bond which is purely chemical by the sense that electrons are shared between two or more nuclie. Also bonding and anti-bonding orbital concept is seen only in the case of covalent bond but not in metallic.
Neither. The metallic "bond" is not really a bond in the way that you are used to thinking of covalent or ionic bonds.scott_alexsk said:Gokul, so what kind of bonds do the atoms form per say, being sigma bonds or pi bonds?
Again, the answer is "neither". You no longer have a concept of a bunch of valence electrons that are localized within a small region of space (known as a bond). The valence electrons are completely delocalized over the entire chunk of the metal. The metallic "bond" is nothing but this sea of electrons.So do the atoms dissociate their electrons to obtain a half filled orbital configuration, a nobel gas configuration, or both in some cases?
This is correct. But I would use the word "overlap" rather than "exceed".Correct me if I am wrong but these bands, control whether or not an element has metal or nonmetal characteristics. Does the location of these bands cause the metalalloy steps (in the periodic table), since over time the conduction band exceeds the valence band? Tell me if I am completely wrong because I would not be surprized if I was.
scott_alexsk said:What is the difference between the valence band and the conductive band? Is it just that the valence band contains electrons and the conductive empty space? Just to clarify.
scott_alexsk said:Now I am a little confused about the formation of conduction bands and valence bands. Are they only created when one has overlapping orbitals? It seems like it is otherwise.
A metallic bond is a type of chemical bond that forms between metal atoms. It is characterized by the sharing of delocalized electrons between the atoms, resulting in a strong attraction between the positive metal ions and the negatively charged electrons.
A metallic bond is formed through a process called metallic bonding, in which the outer electrons of metal atoms become delocalized and are free to move throughout the metal lattice. This allows for the formation of a strong electrostatic attraction between the metal ions and the delocalized electrons.
Metallic bonds are responsible for the unique properties of metals, such as high electrical and thermal conductivity, malleability, and ductility. They also have high melting and boiling points, as well as good strength and hardness.
Metallic bonds differ from other types of chemical bonds, such as covalent and ionic bonds, in that they involve the sharing of delocalized electrons rather than the sharing or transfer of specific electrons between atoms. This results in a more uniform distribution of electrons and a stronger bond.
The strength of a metallic bond is affected by the number of delocalized electrons, the size of the metal ions, and the distance between the metal ions in the metal lattice. Generally, the more delocalized electrons and the smaller the metal ions, the stronger the metallic bond will be.