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Metals - Experiment 13 Revision 2.0 Bonding in Metals and...

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Experiment 13 Revision 2.0 Bonding in Metals and Semiconductors Learn about Metallic Bonding. Learn about Insulators, Semiconductors and Metals. Learn about Hydrogen Insertion. Learn about Light Emitting Diodes. In this laboratory, we will perform a few short exercises to demonstrate the nature of the bonding in metals and semi-conductors. First, we will examine the color of the photons emitted by Gallium Arsenide Phosphide Light Emitting Diodes (LEDs). We will also measure the voltage required to induce a minimum current to flow in these semiconducting materials. These observations will provide us with a feel for how the composition of these diodes influences their Band-Gap energies. Next, we will observe the change induced when these diodes are cooled to liquid Nitrogen temperatures. This will allow us to observe how structural changes in the material also influence the Band-Gap energy. Finally, we will insert Hydrogen atoms into a Tungsten Trioxide (WO 3 ) solid matrix, to prepare H x WO 3 , and observe the resulting change in electrical conductivity of the material. Chemical bonding in Metals is distinctly different than that in Ionically or Covalently Bonded compounds. In Ionically Bonded compounds, such as Sodium Chloride (NaCl), valence electrons are transferred from one atom to another, allowing each atom to obtain a Noble Gas configuration, with the resulting ions electrostatically attracted, or “bonded”, to each other. In Covalently Bonded compounds, such as Methane (CH 4 ), valence electrons of the bonding atoms are shared such that that each achieves a Noble Gas configuration. However, in metals, such as Sodium (Na), there are insufficient valence electrons to be shared to complete each atom’s octet. And, valence electrons will not be transferred from one atom to another to form ions because the atoms have identical electronegativities.
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Instead, the bonding is collective; each atom contributes its valence electrons to form a “sea” of electrons which are shared by all the atoms of the metal. The resulting metal cation cores are held together by their attraction to this “sea” of delocalized electrons. For a concrete example, consider the bonding between two Sodium atoms that form the hypothetical Na 2 molecule. Each atom involved in the bonding has the following electron configuration: Na = 1s 2 2s 2 2p 6 3s 1 = [Ne] 3s 1 Thus, each Na atom has a single valence electron in the 3s orbital. When the two Na atoms approach each other, the 3s valence level orbitals will overlap, creating a region of high electron density where the two valence electrons can be shared by both atoms. The resulting “Molecular Orbital” gives rise to an energetically favorable configuration and is referred to as a Bonding MO.
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However, this picture is a bit simplistic. We have to admit that orbitals are merely representations of Wavefunctions, mathematical waves, which, like other wave pulses, can have a positive or negative Amplitude.
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